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Chemistry – STAR Test Review – Day 1 Ch2 – Classifying Matter

Chemistry – STAR Test Review – Day 1 Ch2 – Classifying Matter 1 . Classify these as elements, compounds, or mixtures. If mixture, decide if heterogeneous or homogeneous. a . silver b . pine tree c . orange juice d . carbon dioxide e . air

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Chemistry – STAR Test Review – Day 1 Ch2 – Classifying Matter

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  1. Chemistry – STAR Test Review – Day 1 Ch2 – Classifying Matter 1. Classify these as elements, compounds, or mixtures. If mixture, decide if heterogeneous or homogeneous. a. silver b. pine tree c. orange juice d. carbon dioxide e. air 2. Classify each of the following changes as physical or chemical. a. bread is baked b. salt dissolves in water c. milk spoils d. A snowflake melts

  2. Ch4,25 – The Atom and Nuclear Processes 1. What is the charge, positive or negative, of the nucleus of every atom? 2. Determine the number of neutrons in each atom. a. nitrogen-15 b. radium-226 3. List the number of protons, neutrons, and electrons in each of the following atoms. a. 2713Al b. 4420Ca Ch5 – The Electron ___ Dalton a. planetary model of the atom, electrons move around the nucleus like planets around sun ___ Bohr b. gold foil experiment – atoms have a dense core called nucleus ___ Rutherford c. father of the modern atomic theory, everything made of atoms 1. Find the Battleship Position, the Electron Configuration, and the Orbital Filling Diagram for these atomic numbers: a. Atomic number 15 b. Atomic number 56 2. What is the frequency of radiation with a wavelength of 5.0 x 10-8 m? 3. What is the energy of a photon whose frequency is 3.0 x 1012 Hz? Ch6 – Periodicity 1. Arrange these elements in order of decreasing atomic size: sulfur, chlorine, aluminum and sodium. Does your arrangement demonstrate a periodic or group trend? 2. Distinguish between the first and second ionization energies of an atom. 3. Chlorine, selenium, and bromine are located near each other on the period table. Which of these elements is the smallest atom? 4. Which particle has the largest radius in each atom /ion pair? a. Na, Na+ b. S, S-2 5. Why don’t the noble gases appear in an eletronegativity table? Day 2 Ch9 – Naming 1. Identify the following elements as metals, metalloids, or nonmetals. a. gold b. silicon 2. The melting point of a compound is 1240’C, and ice melts at 0’C. Is this compound an ionic or molecular compound? 3. Write the formulas for these compounds. a. barium fluoride b. calcium carbonate c. lithium hypochlorite d. tin(II) hydroxide e. tetraiodinenonoxide 4. Write names for these compounds. a. Al(OH)3 b. NaI c. Sn3(PO4)2 d. CS2 e. Cl2O7 f. N2O5 Ch10 – The Mole 1. How many moles are 2.80x1024 atoms of silicon? 2. Find the gram molecular mass of C3H7OH. 3. What is the mass of 2.11 x 1024 molecules of sulfur dioxide? 4. Find the volume at STP of 3.20 x 10-3 mol CO2 gas. 5. At STP, what volume does 12.2 g of fluorine gas occupy? 6. 13.0 x 1024 molecules of argon occupy what volume at STP? Ch11 – Chemical Reactions 1. Balance the following equations: a. ___ PbO2 ___ PbO + ___ O2 b. ___ Li + ___ FeBr2 ___ LiBr + ___ Fe c. ___ Al + ___ CuSO4 ___ Al2(SO4)3 + ___ Cu 2. Rewrite these word equations as balanced chemical equations. a. iron (III) chloride + calcium hydroxide  iron (III) hydroxide + calcium chloride b. sodium + water  sodium hydroxide + hydrogen 3. Write the balanced chemical equation for the following combination reaction: Mg + O2 4. Write the balanced chemical equation for the following decomposition reaction: HBr 5. Write the balanced chemical equation for the following single-replacement reaction: Ag(s) + KNO3(aq) 6. Write a balanced chemical equation for the following double replacement reactions: HCl(aq) + Ca(OH)2(aq) 7. After balancing the equation, identify the type. a. ___ C2H6 + ___ O2 ___ CO2 + ___ H2O b. ___ Pb(NO3)2 + ___ NaI ___ PbI2 + ___ NaNO3 Day 3 Ch12 – Stoichiometry 1. Determine the mass of lithium hydroxide produced when 0.38 g of lithium nitride reacts with water according to the following equation: Li3N + 3H2O → NH3 + 3LiOH 2. Determine the mass of carbon dioxide produced when 0.85 g of butane reacts with oxygen according to the following equation: 2C4H10 + 13O2 → 8CO2 + 10H2O 3. ___ C2H5OH + ___ O2  ___ CO2 + ___ H2O Given 26 L O2, at STP, how many liters of CO2 will be produced? 4. Carbon disulfide is an important industrial solvent. It is prepared by the reaction of coke with sulfur dioxide. 5C + 2SO2 CS2 + 4CO how many moles of CS2 form when 2.7 mol of C reacts? 5. The reaction between fluorine with ammonia produces dinitrogentetrafluoride and hydrogen fluoride. 5F2(g) + 2NH3(g)  N2F4(g) + 6HF(g) if you have 66.6 g of NH3, how many liters of F2 are required for complete reaction? Ch13,14 – Gas Laws 1. According to the kinetic theory, gases consist of particles that: a. occupy considerable volume b. are close together c. exert attractive and repulsive forces among themselves d. move rapidly in constant random motion e. None of these 2. At sea level, atmospheric pressure is equal to __________ 3. STP refers to ________________________ 4. Absolute zero is: a. the temperature at which the motion of particles theoretically ceases b. 0 K c. a temperature that has never been produced in the laboratory d. All of these e. None of these 5. At 80 K, the particles of gas have: a. twice the average kinetic energy of the same particles at 40K b. half the average kinetic energy of the same particles at 40K c. one-forth the average kinetic energy of the same particles at 20K d. None of these 6. Compare H2O as a gas, liquid, and solid. a. Which state has the most kinetic energy? b. Which state has a definite shape? c. Which state has its volume most affected by pressure? 7. Water could be made to boil at 92ºC by: (think about pressure) 8. As the temperature of a fixed volume of a gas increases, the pressure will __________ 9. The volume of a gas is doubled while the temperature is held constant. The pressure of the gas __________ 10. If a sample of oxygen occupies a volume of 2.15 L at a pressure of 58.0 kPa and a temperature of 25C, what volume would this sample occupy at 101.3 kPa and 0C? (Combined Gas Law) 11. The volume (in L) that would be occupied by 5.00 mol of O2 at STP is: (PV = nRT) Day 4 Chapter 7,8 – Ionic and Covalent Bonds 1. How many valence electrons does an atom of any element in Group 15 have? 2. Draw the electron dot structure for an atom of fluorine. 3. The general electron dot structure . X : could represent which column on the periodic table? 4. When a magnesium atom loses its valence electrons, what is the charge on the resulting ion? 5. What is the electron configuration of oxygen and also the oxide ion, O -2? 6. An ionic compound is: a. generally a salt b. held together by ionic bonds c. composed of anions and cations d. all of the above 7. Which of these is not a characteristic of most ionic compounds? a. solid at room temperature b. conducts an electric current when melted c. has a low melting point d. produced by reaction between metallic and nonmetallic elements 8. A covalent bond forms? a. when as element becomes a noble gas b. when atoms share electrons c. between metals and nonmetals d. when electrons are transferred from one atom to another 9. Name one element with eight valence electrons. 10. In general metals react by: a. losing valence electrons b. gaining valence electrons c. sharing valence electrons d. sometimes gaining and sometimes losing valence electrons. 11. An ion of K has the same electron configuration as what noble gas? 12. Will the following atoms combine to form ionic bonds, polar covalent bonds, or nonpolar covalent bonds? a. sodium and bromine b. nitrogen and oxygen c. hydrogen and carbon Chapter 15,16 – Properties of Water and Solutions 1. The high surface tension of water is due to ________________________ 2. Is a solution is a heterogeneous or homogeneous mixture? Can the solute be filtered out of the solvent? What does it mean to say a solution is saturated? 3. What is an electrolyte? 4. How many water molecules are attached in the substance copper sulfate pentahydrate? 5. Which of these would you expect to be soluble in the nonpolar solvent carbon disulfide, CS2? a. SnS2 b. CaCO3 c. CCl4 d. H2O 6. Which of the following would be expected to dissolve very readily in water? a. CH4 b. H2 c. NaOH d. CCl4 7. What 3 things can you do to usually make a substance dissolve faster in a solvent? 8. To increase the solubility of a gas at constant temperature, would it be better to increase or decrease the pressure above the liquid? 9. What is the molarity of a 200 mL solution in which 0.2 mole of sodium bromide is dissolved? Day 5 Chapter 17 – Thermochemistry 1. Water has a specific heat capacity of 4.18 J/(g.ºC). Iron has a specific heat capacity of 0.44 J/(g.ºC). Which can absorb more heat energy without changing its temperature as much? 2. The boiling point of an unknown solution is 125ºC. What is this temperature in Kelvin? 3. A beaker contains 250 mL of water and is placed on a hot plate. The water is heated. Heat energy comes from the hot plate, enters the water, and the water’s temperature rises until it starts to boil. When the water starts to boil, the temperature stops rising, yet the hot plate continues to pump heat energy into the water. How can this be? 4. Given the reaction: N2(g) + 3H2(g) 2 NH3(g) ΔH = – 92 kJ Is the reaction endothermic or exothermic? Re-write the reaction to include the heat. If 5 L of Hydrogen react with excess nitrogen, how much heat is produced? Chapter 18 – Reaction Rates and Equilibrium 1. Based on the reaction given in #4 above, if someone keeps adding more and more hydrogen gas to an excess of nitrogen, what will happen to the amount of ammonia produced? 2. If the reaction in #4 above took place in a large balloon and was allowed to reach equilibrium, what would happen if someone squeezed the balloon into a smaller volume? 3. Which has more entropy, solid iodine or iodine that has sublimed into a gas? 4. At 25ºC, the following reaction occurs: 4NH3(g) + 7O2(g) 4NO2(g) + 6H2O(g) + 1130.6 kJ a. The change in entropy, ΔS = – 110.8 J/K. Does this reaction have a favorable change in entropy? b. Is this reaction spontaneous, based on the Gibb’s Free Energy Equation? Chemistry – The Missing Fourth Quarter Stuff Chapter 19 – Acids and Bases 1. Which of the following substances are acids and which are bases? While you’re at it, why don’t you name them as well! a. HCl b. H2SO4 c. NaOH d. Ca(OH)2 e. NH3 f. HClO3 g. HClO 2. The pH of pure water is 7. If you squirt some lemon juice in it, its pH drops to 6. Is it considered acidic or basic? 3. The pH of a concentrated HCl solution is 1. The pH of the lemon juice is 6. Please write something about the relationship between the pH and the strength of the acid. 4. Bases have pH values above 7. Ammonia has a pH of 8. A NaOH solution has a pH of 14. Which would neutralize HCl solution better? Chapter 20 – Neutralizations 1. Sodium hydroxide neutralizes hydrochloric acid. a. What are the products? b. If your stomach contains 500 mLs of a 0.1 M HCl solution, how many grams of NaOH could you swallow to neutralize this acid? (Please don’t attempt this!) Carbon Chains and Functional Groups 1. Octane, commonly found in gasoline, is an alkane, a chain of carbon atoms, one of many hydrocarbons we have mentioned this year. a. Please draw a model of the molecule. b. Explain why the molecule is saturated. 2. Ethene, C2H4, is a non-saturated hydrocarbon. Show how it can become saturated if it forms the alkane named ethane, C2H6. 3. Butane, C4H10, is an alkane with 2 isomers. Draw them. 4. Ethanol, C2H5OH, is an alcohol group. Show how it is different from ethane.

  3. 4. Tell how alpha, beta and gamma radiation are distinguished on the basis of the following: a. mass b. charge c. penetration power 5. Complete these nuclear reactions: a. 3015P  _____ + 0-1e b. _____ 147N + 0-1e

  4. 6. Write nuclear reaction for the beta decay of the following isotope: 9038Sr 7. Write the nuclear reaction for this word equation: a. Radon-222 emits an alpha particle to form polonium-218.

  5. Ch5 – The Electron ___ Dalton a. planetary model of the atom, electrons move around the nucleus like planets around sun   ___ Bohr b. gold foil experiment – atoms have a dense core called nucleus   ___ Rutherford c. father of the modern atomic theory, everything made of atoms 1. Find the Battleship Position, the Electron Configuration, and the Orbital Filling Diagram for these atomic numbers: a. Atomic number 15 b. Atomic number 56

  6. 2. What is the frequency of radiation with a wavelength of 5.0 x 10-8m? 3. What is the energy of a photon whose frequency is 3.0 x 1012Hz? Ch6 – Periodicity 1. Arrange these elements in order of decreasing atomic size: sulfur, chlorine, aluminum and sodium. Does your arrangement demonstrate a periodic or group trend?

  7. 2. Distinguish between the first and second ionization energies of an atom. 3. Chlorine, selenium, and bromine are located near each other on the period table. Which of these elements is the smallest atom?

  8. 4. Which particle has the largest radius in each atom /ion pair? a. Na, Na+ b. S, S-2 5. Why don’t the noble gases appear in an eletronegativity table?

  9. Day 2 Ch9 – Naming 1. Identify the following elements as metals, metalloids, or nonmetals. a. gold b. silicon 2. The melting point of a compound is 1240’C, and ice melts at 0’C. Is this compound an ionic or molecular compound?

  10. 3. Write the formulas for these compounds. a. barium fluoride b. calcium carbonate c. lithium hypochlorite d. tin(II) hydroxide e. tetraiodinenonoxide

  11. 4. Write names for these compounds. a. Al(OH)3 b. NaI c. Sn3(PO4)2 d. CS2 e. Cl2O7 f. N2O5

  12. Ch10 – The Mole 1. How many moles are 2.80x1024 atoms of silicon? 2. Find the gram molecular mass of C3H7OH. 3. What is the mass of 2.11 x 1024 molecules of sulfur dioxide?

  13. 4. Find the volume at STP of 3.20 x 10-3 mol CO2 gas. 5. At STP, what volume does 12.2 g of fluorine gas occupy? 6. 13.0 x 1024 molecules of argon occupy what volume at STP?

  14. Ch11 – Chemical Reactions 1. Balance the following equations: a. ___ PbO2 ___ PbO + ___ O2 b. ___ Li + ___ FeBr2 ___ LiBr + ___ Fe c. ___ Al + ___ CuSO4 ___ Al2(SO4)3 + ___ Cu

  15. 2. Rewrite these word equations as balanced chemical equations. a. iron (III) chloride + calcium hydroxide  iron (III) hydroxide + calcium chloride b. sodium + water  sodium hydroxide + hydrogen

  16. 3. Write the balanced chemical equation for the following combination reaction: Mg + O2 4. Write the balanced chemical equation for the following decomposition reaction: HBr 5. Write the balanced chemical equation for the following single-replacement reaction: Ag(s) + KNO3(aq)

  17. 6. Write a balanced chemical equation for the following double replacement reactions: HCl(aq) + Ca(OH)2(aq) 7. After balancing the equation, identify the type. a. ___ C2H6 + ___ O2 ___ CO2 + ___ H2O b. ___ Pb(NO3)2 + ___ NaI ___ PbI2 + ___ NaNO3

  18. Day 3 Ch12 – Stoichiometry 1. Determine the mass of lithium hydroxide produced when 0.38 g of lithium nitride reacts with water according to the following equation: Li3N + 3H2O → NH3 + 3LiOH

  19. 2. Determine the mass of carbon dioxide produced when 0.85 g of butane reacts with oxygen according to the following equation: 2C4H10 + 13O2 → 8CO2 + 10H2O  3. ___ C2H5OH + ___ O2  ___ CO2 + ___ H2O Given 26 L O2, at STP, how many liters of CO2 will be produced?

  20. 4. Carbon disulfide is an important industrial solvent. It is prepared by the reaction of coke with sulfur dioxide. 5C + 2SO2 CS2 + 4CO how many moles of CS2 form when 2.7 mol of C reacts?  5. The reaction between fluorine with ammonia produces dinitrogentetrafluoride and hydrogen fluoride. 5F2(g) + 2NH3(g)  N2F4(g) + 6HF(g) if you have 66.6 g of NH3, how many liters of F2 are required for complete reaction?

  21. Ch13,14 – Gas Laws 1. According to the kinetic theory, gases consist of particles that: a. occupy considerable volume b. are close together c. exert attractive and repulsive forces among themselves d. move rapidly in constant random motion e. None of these 2. At sea level, atmospheric pressure is equal to __________ 3. STP refers to ________________________ 4. Absolute zero is: a. the temperature at which the motion of particles theoretically ceases b. 0 K c. a temperature that has never been produced in the laboratory d. All of these e. None of these

  22. 5. At 80 K, the particles of gas have: a. twice the average kinetic energy of the same particles at 40K b. half the average kinetic energy of the same particles at 40K c. one-forth the average kinetic energy of the same particles at 20K d. None of these  6. Compare H2O as a gas, liquid, and solid. a. Which state has the most kinetic energy? b. Which state has a definite shape? c. Which state has its volume most affected by pressure? 7. Water could be made to boil at 92ºC by: (think about pressure)

  23. 8. As the temperature of a fixed volume of a gas increases, the pressure will __________ 9. The volume of a gas is doubled while the temperature is held constant. The pressure of the gas __________ 10. If a sample of oxygen occupies a volume of 2.15 L at a pressure of 58.0 kPa and a temperature of 25C, what volume would this sample occupy at 101.3 kPa and 0C? (Combined Gas Law)

  24. 11. The volume (in L) that would be occupied by 5.00 mol of O2 at STP is: (PV = nRT)

  25. Day 4 Chapter 7,8 – Ionic and Covalent Bonds 1. How many valence electrons does an atom of any element in Group 15 have? 2. Draw the electron dot structure for an atom of fluorine. 3. The general electron dot structure . X : could represent which column on the periodic table? 4. When a magnesium atom loses its valence electrons, what is the charge on the resulting ion?

  26. 5. What is the electron configuration of oxygen and also the oxide ion, O -2? 6. An ionic compound is: a. generally a salt b. held together by ionic bonds c. composed of anions and cations d. all of the above 7. Which of these is not a characteristic of most ionic compounds? a. solid at room temperature b. conducts an electric current when melted c. has a low melting point d. produced by reaction between metallic and nonmetallic elements

  27. 8. A covalent bond forms? a. when as element becomes a noble gas b. when atoms share electrons c. between metals and nonmetals d. when electrons are transferred from one atom to another 9. Name one element with eight valence electrons. 10. In general metals react by: a. losing valence electrons b. gaining valence electrons c. sharing valence electrons d. sometimes gaining and sometimes losing valence electrons. 11. An ion of K has the same electron configuration as what noble gas?

  28. 12. Will the following atoms combine to form ionic bonds, polar covalent bonds, or nonpolar covalent bonds? a. sodium and bromine b. nitrogen and oxygen c. hydrogen and carbon Chapter 15,16 – Properties of Water and Solutions 1. The high surface tension of water is due to ___________________ 2. Is a solution is a heterogeneous or homogeneous mixture? Can the solute be filtered out of the solvent? What does it mean to say a solution is saturated? 3. What is an electrolyte?

  29. 4. How many water molecules are attached in the substance copper sulfate pentahydrate? 5. Which of these would you expect to be soluble in the nonpolar solvent carbon disulfide, CS2? a. SnS2 b. CaCO3 c. CCl4 d. H2O 6. Which of the following would be expected to dissolve very readily in water? a. CH4 b. H2 c. NaOH d. CCl4

  30. 7. What 3 things can you do to usually make a substance dissolve faster in a solvent? 8. To increase the solubility of a gas at constant temperature, would it be better to increase or decrease the pressure above the liquid? 9. What is the molarity of a 200 mL solution in which 0.2 mole of sodium bromide is dissolved?

  31. Day 5 Chapter 17 – Thermochemistry 1. Water has a specific heat capacity of 4.18 J/(g.ºC). Iron has a specific heat capacity of 0.44 J/(g.ºC). Which can absorb more heat energy without changing its temp as much? 2. The boiling point of an unknown solution is 125ºC. What is this temperature in Kelvin?

  32. 3. A beaker contains 250 mL of water and is placed on a hot plate. The water is heated. Heat energy comes from the hot plate, enters the water, and the water’s temperature rises until it starts to boil. When the water starts to boil, the temperature stops rising, yet the hot plate continues to pump heat energy into the water. How can this be? 4. Given the reaction: N2(g) + 3H2(g) 2 NH3(g) ΔH = – 92 kJ Is the reaction endothermic or exothermic? Re-write the reaction to include the heat. If 5 L of Hydrogen react with excess nitrogen, how much heat is produced?

  33. Chapter 18 – Reaction Rates and Equilibrium 1. N2(g) + 3H2(g) 2 NH3(g), if someone keeps adding more and more hydrogen gas to an excess of nitrogen, what will happen to the amount of ammonia produced? 2. If that reaction took place in a large balloon and was allowed to reach equilibrium, what would happen if someone squeezed the balloon into a smaller volume? 3. Which has more entropy, solid iodine or iodine that has sublimed into a gas?

  34. 4. At 25ºC, the following reaction occurs: 4NH3(g) + 7O2(g) 4NO2(g) + 6H2O(g) + 1130.6 kJ a. The change in entropy, ΔS = – 110.8 J/K. Does this reaction have a favorable change in entropy? b. Is this reaction spontaneous, based on the Gibb’s Free Energy Equation?

  35. Chemistry – The Missing Fourth Quarter Stuff Chapter 19 – Acids and Bases 1. Which of the following substances are acids and which are bases? While you’re at it, why don’t you name them as well! a. HCl b. H2SO4 c. NaOH d. Ca(OH)2 e. NH3 f. HClO3 g. HClO 2. The pH of pure water is 7. If you squirt some lemon juice in it, its pH drops to 6. Is it considered acidic or basic?

  36. 3. The pH of a concentrated HCl solution is 1. The pH of the lemon juice is 6. Please write something about the relationship between the pH and the strength of the acid. 4. Bases have pH values above 7. Ammonia has a pH of 8. A NaOH solution has a pH of 14. Which would neutralize HCl solution better?

  37. Chapter 20 – Neutralizations 1. Sodium hydroxide neutralizes hydrochloric acid. a. What are the products? b. If your stomach contains 500 mLs of a 0.1 M HCl solution, how many grams of NaOH could you swallow to neutralize this acid? (Please don’t attempt this!)

  38. Carbon Chains and Functional Groups 1. Octane, commonly found in gasoline, is an alkane, a chain of carbon atoms, one of many hydrocarbons we have mentioned this year. a. Please draw a model of the molecule. b. Explain why the molecule is saturated. 2. Ethene, C2H4, is a non-saturated hydrocarbon. Show how it can become saturated if it forms the alkane named ethane, C2H6.

  39. 3. Butane, C4H10, is an alkane with 2 isomers. Draw them. 4. Ethanol, C2H5OH, is an alcohol group. Show how it is different from ethane.

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