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Scientific Measurement Place a checkmark next to each item that you can do. If a sample problem is given, complete it as evidence. _____1. I can calculate percent error using the equation on Table T. A researcher measures the mass of a sample to be 5.51 g. The actual mass of the sample is known to be 5.80 g. Calculate the error of the measurement. % error = 5.51-5.80 x 100 = 5% (can make the answer positive) 5.80 How many joules are in 9.3 kJ? 9,300 J _____2. I can convert between different metric units. How many grams are in 39,983 mg? 39.983 g How many milliliters are in 452 L? 452,000 mL What is the volume of 54.6 g of beryllium (density = 1.85 g/cm3)? _____3. I can solve for any variable of the density equation. Volume = 54.6g / 1.85 g/cm3 = 29.5 cm3 What is the density of a 27.8 g substance with a volume of 2.3 mL? Density = 27.8 g / 2.3 mL = 12 g/mL _____4. I can determine the number of significant figures in a measurement. _____5. I can determine the answer to a math problem to the correct number of significant figures. How many significant figures are there in 30.50 cm? 4 How many significant figures are there in 400 sec? 1 To the correct number of significant figures, what is the answer to 5.93 mL + 4.6 mL? 10.5 mL To the correct number of significant figures, what is the answer to 5.93 cm * 4.6 cm? 27 cm2 _____6. I can read the meniscus on a graduated cylinder to the correct number of significant figures. The volume is 75.7 mL. (only 1 decimal) (The 7 can be any one estimated number) _____7. I can convert numbers into scientific notation from standard notation. Convert 87,394,000,000,000 to scientific notation. 8.7394 x 1013 _____8. I can convert numbers into standard notation from scientific notation. Convert 5.8 x 10 to standard notation. 5,800,000,000 Convert 0.0000040934 to scientific notation. 4.0934 x 10-6 9 -5 Convert 4.3 x 10 to standard notation. 0.000043 Unit 1: Matter & Energy Place a checkmark next to each item that you can do. If a sample problem is given, complete it as evidence. Definitions: Atom: the smallest particle of matter that retains the properties of an element. _____1. I can define the terms atom, element, compound, and mixture. Element: a substance that cannot be broken down; made of one type of atom. Compound: atoms of two or more different elements chemically combined in a fixed (definite) ratio. Mixture: two or more substances physically combined in a variable ratio. Put each of the following examples into the correct column. _____2. I can classify substances as a pure substance (element or compound) or a mixture. _____3. I can list the 7 diatomic elements. _____4. I can identify a binary compound. _____5. I can define homogeneous mixture and heterogeneous mixture. Examples: C12H22O11, NaCl, Fe, salt water, air, CO2, H2, Ar, soda Element Compound Mixture Fe, H2, Ar C12H22O11, NaCl, CO2 The seven diatomic elements are: Br2 I2 N2 Cl2 H2 O2 F2 Which of the following is a binary compound? A) NaOH B) H3PO4 C) H2O D) MgSO4 Definitions: Homogeneous Mixture: two or more substances physically combined with uniform composition; no visible differences. Heterogeneous Mixture: two or more substances physically combined with a non-uniform composition; visibly different parts. Two examples of Homogeneous Mixtures: A) salt water (any solution – aq) _____6. I can give examples of homogeneous and heterogeneous mixtures. salt water, air, soda B) Kool Aid Two examples of Heterogeneous Mixtures: A) soil B) sand in water _____7. I can determine how matter will be separated using filtration. _____8. I can describe how matter can be separated using distillation. When a mixture of sand, salt, sugar, and water is filtered, what passes through the filter? salt, sugar, water Which physical property makes it possible to separate the components of crude oil by means of distillation? difference in boiling points To separate a mixture of salt and water, use evaporation or distillation. _____9. I can state which separation process is best for a given situation. To separate a mixture of ethanol and water, use distillation. To separate a mixture of sand and water, use filtration. Monoatomic Element Diatomic Element (same colors) ____10. I can draw particle Compound diagrams to represent a monoatomic element, a diatomic element, a compound, a mixture of 2 compounds, and a mixture (different colors) of an element and a compound. Mixture of an element and compound Mixture of 2 compounds All chemical reactions have conservation of _____11. I can answer questions about the Law of Conservation of Matter (Mass). A) mass only C) mass and charge only B) charge and energy only D) mass, charge, and energy Given the reaction: C6H12O6 + 6 O2 6 CO2 + 6 H2O, determine the mass of CO2 produced when 9 grams of glucose completely reacts with 9.6 grams of oxygen to produce 5.4 grams of water. 13.2 g (total mass of reactants = total mass of products) Write “P” for physical or “C” for chemical on the line provided. P Copper (II) sulfate is blue. C Copper reacts with oxygen. _____12. I can classify a P Copper can be made into wire. property as physical or chemical. P Copper has a density of 8.96 g/cm3. P Copper melts at 1358 K. P Copper doesn’t dissolve in water. Write “P” for physical or “C” for chemical on the line provided. P Copper (II) sulfate dissolves in water. C Copper reacts with oxygen to form solid copper (I) oxide. _____13. I can classify a change as physical or chemical. P Solid copper is melted. P A chunk of copper is pounded flat. P Copper and zinc are mixed to form brass. P A large piece of copper is chopped in half. C Copper reacts with bromine to form copper (II) bromide. Circle the particle diagram that best represents Substance A after a physical change has occurred. _____14. In a particle diagram, I can distinguish between a physical change and a chemical change. Substance A Draw a particle diagram to represent atoms of Li in each phase. Solid Liquid Gas ____15. I can use particle diagrams to show the arrangement and spacing of atoms/molecules in different phases. _____16. I can compare solids, liquids, and gases in terms of their shape, volume, particle distance, and attraction between particles. Shape Volume Particle Distance Attraction Between Particles Solid Definite Definite Close together Liquid Indefinite Definite Random, but close together Gas Indefinite Indefinite Random and far apart Strong Weak Very Strong *Gases fill the container most completely Fusion: change from solid to liquid. Solidification: change from liquid to solid. Condensation: change from gas to liquid. Vaporization: change from liquid to gas. _____17. I can state the change of phase occurring during each phase change. Melting: change from solid to liquid. Boiling: change from liquid to gas. Sublimation: change from solid to gas. Deposition: change from gas to solid. Freezing: change from liquid to solid. _____18. I can answer questions about the Law of Conservation of Energy. As electrical energy is converted into heat energy, the total amount of energy A) decreases B) increases C) remains the same Definitions: Kinetic Energy: energy a substance has due to its motion (related to temperature). Potential Energy: energy a substance has stored in its bonds (related to distance between particles). _____19. I can define kinetic energy, potential energy, temperature, heat, endothermic, and exothermic. Temperature: a measure of the average kinetic energy of a substance. Heat: energy transferred from an object of high temperature to an object of low temperature. Endothermic: a process in which energy is absorbed. Exothermic: a process in which energy is released. Indicate whether the change is exothermic or endothermic: fusion/melting endothermic solidification/freezing exothermic _____20. I can indicate if a phase change is exothermic or endothermic. condensation exothermic vaporization/boiling endothermic sublimation endothermic deposition exothermic Which sample of water contains particles having the highest kinetic energy? _____21. I can answer questions about average kinetic energy. A) 25 mL of water at 95˚C B) 45 mL of water at 75˚C C) 75 mL of water at 75˚C D) 95 mL of water at 25˚C Object A at 40˚C and object B at 80˚C are placed in contact with each other. Which statement describes the heat flow between objects? _____22. I can state the direction A) Heat flows from object A to object B B) Heat flows from object B to object A of heat flow. C) Heat flows in both directions between the objects D) No heat flow occurs between the objects _____23. I can convert between Celsius and Kelvin. _____24. Given a heating/cooling curve, I can identify the process occurring during each segment. What Kelvin temperature is equal to 200˚C? 473 K What Celsius temperature is equal to 200 K? -73˚C AB = heating or cooling a solid BC = melting or freezing CD = heating or cooling a liquid DE = vaporizing/boiling or condensing EF = heating or cooling a gas On the graph, circle the sections that show a change in potential energy. _____25. Given a heating/cooling curve, I can determine which sections of the curve show changes in potential energy. On the graph, circle the sections that show a change in kinetic energy. _____26. Given a heating/cooling curve, I can determine which sections of the curve show changes in kinetic energy. On the graph, circle the sections that show a phase change. _____27. Given a heating/cooling curve, I can determine which sections of the curve show phase changes. 113oC _____28. Given a heating/cooling curve, I can determine the temperature at which a substance melts/freezes or vaporizes/condenses. 53oC What is the freezing point of this substance? 53˚C What is the boiling point of this substance? 113˚C _____29. I can state the temperature at which water freezes in ˚C and K. _____30. I can state the temperature at which water melts in ˚C and K. _____31. I can state the temperature at which water boils in ˚C and K. _____32. I can state the temperature at which water condenses in ˚C and K. _____33. I can state the unit used to measure energy. What is the freezing point of water in ˚C and K? 0˚C or 273 K What is the melting point of water in ˚C and K? 0˚C or 273 K What is the boiling point of water in ˚C and K? 100˚C or 373 K What is the condensing point of water in ˚C and K? 100˚C or 373 K Energy is measured in joules. Definitions: Specific Heat Capacity: amount of heat required to increase the temperature of 1 gram of a substance 1˚C (or 1 K). _____34. I can define specific heat capacity, heat of fusion, and heat of vaporization. Heat of Fusion: amount of heat that needs to be absorbed to melt 1 gram of a substance (or released to freeze 1 gram of a substance). Heat of Vaporization: amount of heat that needs to be absorbed to vaporize 1 gram of a substance (or released to condense 1 gram of a substance). How much heat is needed to vaporize 10 g of water at 100˚C? q = mHV q = 10(2,260) = 22,600 J (Hv on Table B, don’t plug in temp.) How much heat is needed to change the temperature of 100.0 g of water from 20˚C to 35˚C? q = mC∆T q = 100(4.18)(15) = 6,270 J (C on Table B) _____35. I can use the heat equations to solve for any variable. How much heat is needed to melt 20 g of ice at 0˚C? q = mHf q = 20(334) = 6,680 J (Hf on Table B, don’t plug in temp.) How many grams of water can be heated by 15˚C using 13,500 J of heat? q = mC∆T 13,500 = m(4.18)(15) = 215 g It takes 5,210 J of heat to melt 50 g of ethanol at its melting point. What is the heat of fusion of ethanol? q = mHf 5,210 = 50(Hf) = 104.2 J/g (solving for Hf, substance is not water so you cannot use Table B value) The five parts of the Kinetic Molecular Theory are: A. Particles of an ideal gas move in constant, random, straight line motion. _____36. I can state the 5 parts of the Kinetic Molecular Theory. B. Particles of an ideal gas have elastic collisions (energy is not lost). C. The size of the particles is so small compared to the space between the particles that the volume of the actual gas particles is negligible. D. Particles of an ideal gas have no forces of attraction between them. E. Particles of an ideal gas collide with container walls causing pressure. _____37. I can define an ideal gas. Definition: Ideal Gas: follows all of the parts of the Kinetic Molecular Theory. _____38. I can state why real gases do not behave like ideal gases. Real gases do not behave like ideal gases because real gases have some _____39. I can state the two elements that act ideally most of the time. _____40. I can state the conditions under which real gases behave most like ideal gases. The two elements that act ideally most of the time are hydrogen & volume and some attraction. helium. A real gas will act most ideally under the conditions of low pressure and high temperature. (PLIGHT) STP stands for Standard Temperature and Pressure. _____41. I can state the meaning of “STP” and the Table on which it can be found. The values can be found on Table A. Standard Pressure = 1 atm or 101.3 kPa Standard Temperature = 0˚C or 273 K Avogadro’s Hypothesis says that samples of a gas that are at the same temperature, pressure, and volume will contain the same number of molecules. The actual identity of the substances does not matter. _____42. I can state Avogadro’s Hypothesis (Law) and answer questions based on this concept. Which cylinder contains the same number of molecules at STP as a 2 L cylinder of H2 (g) at STP? A) 1 L cylinder of O2 (g) C) 2 L cylinder of CH4 (g) B) 1.5 L cylinder of NH3 (g) D) 4 L cylinder of He (g) _____43. I can state the relationship between pressure and volume for gases. At constant temperature, as the pressure on a gas increases, the volume _____44. I can state the relationship between temperature and volume for gases. At constant pressure, as the temperature of a gas increases, the volume _____45. I can state the relationship between temperature and pressure for gases. In a fixed container (has constant volume), as the temperature of a gas decreases. increases. increases, the pressure increases. Draw a curve for each of the graphs below: _____46. I can recognize the relationships stated above in graph form. P V V P T T A rigid cylinder with a movable piston contains a 2 L sample of neon gas at STP. What is the volume when its temperature is increased to 30˚C while its pressure is decreased to 90 kPa? P1V1 = P2V2 101.3(2) = 90V2 V2 = 2.5 L T1 T2 273 303 _____47. I can solve problems using the combined gas law. * STP values on Table A * Choose the pressure with the same units as the final pressure * Temperature needs to be in Kelvin * Use the temperature conversion equation on Table T Definitions: Vapor Pressure: the force a gas exerts above a liquid. _____48. I can define vapor pressure, boiling point, and volatile. Boiling Point: the temperature at which the vapor pressure is equal to the atmospheric pressure. Volatile: easily turns into a vapor (easily evaporates). _____49. I can state the conditions of pressure that are used for normal boiling points. The normal boiling point of a substance occurs at a pressure of 1 atm or 101.3 kPa. This is known as standard pressure and can be found on Table A. _____50. I can state the As the temperature of a substance increases, its vapor pressure increases. relationship between temperature and vapor pressure. _____51. I can state the As the atmospheric pressure increases, the boiling point increases. relationship between atmospheric pressure and boiling point. _____52. I can state the As the attraction between molecules increases, the boiling point increases. relationship between molecular attraction and boiling point. _____53. I can determine strength of attraction between molecules using Table H. Based on Table H, which substance has the weakest intermolecular forces? Propanone, because at any temperature it has the highest vapor pressure. This means it is easily converted into a vapor (weak forces). _____54. I can determine the vapor pressure of ethanol, ethanoic acid, propanone, or water at a given temperature. What is the vapor pressure of ethanol at 56˚C? 39 kPa _____55. I can determine the temperature of ethanol, ethanoic acid, propanone, or water at a given vapor pressure. When the vapor pressure of water is 30 kPa, what is the temperature of the water? 70˚C Unit 2: The Atom Place a checkmark next to each item that you can do. If a sample problem is given, complete it as evidence. _____1. I can describe John Dalton’s contribution to our understanding of the atom. Dalton’s Model: the atom is the basic unit of matter; atoms of the same element have the same properties, atoms of different elements have different properties. Diagram: Thomson’s Model: “Plum Pudding” Model – small, negative electrons embedded in a large positive sphere. _____2. I can describe JJ Thomson’s contribution to our understanding of the atom. Diagram: Rutherford’s Experiment: directed positive alpha particles at a thin sheet of gold foil and observed their path. Experiment Results: most alpha particles went straight through, a few were slightly deflected, some bounced back. _____3. I can describe Ernest Rutherford’s contribution to our understanding of the atom. Experiment Conclusions: the atom is mostly empty space, with the mass concentrated in the small, dense, positively charged nucleus. Diagram: Bohr’s Model: “Planetary Model” – electrons are in definite orbits (circular paths) around the nucleus. _____4. I can describe Niels Bohr’s contribution to our understanding of the atom. Diagram: Where are electrons likely to be found according to the modern model? Orbitals – regions of probable location _____5. I can describe the modern model (wave mechanical model) of the atom. Diagram: Particle #1 Particle #2 Particle #3 Name Proton Neutron Electron Charge +1 0 -1 Mass 1 amu 1 amu 0 amu Location in Atom Nucleus Nucleus Outside Nucleus _____6. I can state the three subatomic particles, their location in an atom, and the magnitude of their charges and their masses (in amu). _____7. I can identify the nuclear charge of an atom. _____8. I can explain why atoms are electrically neutral. _____9. I can define atomic number and mass number. What is the charge of an oxygen-18 nucleus? +8 Atoms are electrically neutral because the number of protons is equal to the number of electrons. Definitions: Atomic Number: the number of protons in the nucleus; defines the identity of the element. Mass Number: the sum of the protons and neutrons. _____10. I can use the Periodic Table to determine the atomic number of an element. How many protons are in an atom of selenium? 34 _____11. I can determine the mass number of an atom. What is the mass number of an atom that has 20 protons, 19 neutrons, and 20 electrons? 20 + 19 = 39 _____12. I can represent an atom using different notations. How many protons are in an atom of silicon? 14 What are three other notations that can be used to represent 80Br? 80 Br Bromine-80 Br-80 35 _____13. Given the mass number, I can determine the number of protons, neutrons, and electrons in an atom. In an atom of 212Po, how many protons are present? 84 In an atom of 212Po, how many electrons are present? 84 In an atom of 212Po, how many neutrons are present? 128 Definitions: Ion: charged particle due to the loss or gain of electrons. _____14. I can define ion, cation, and anion. Cation: a positive ion formed when an atom loses one or more electrons. Anion: a negative ion formed when an atom gains one or more electrons. How many protons are in 18O-2? 8 How many neutrons are in 18O-2? 10 _____15. Given the mass number and the charge, I can determine the number of protons, neutrons, and electrons in an ion. How many electrons are in 18O-2? 10 What is this ion’s electron configuration? 2-8 How many electrons are in 27Al+3? 10 What is this ion’s electron configuration? 2-8 _____16. I can define isotope. _____17. I can identify isotopes based on their symbols. _____18. I can calculate average atomic mass given the masses of the naturally occurring isotopes and the percent abundances. Definition: Isotope: atoms of the same element (same atomic #, same # of protons) with a different mass # (different # of neutrons). Which symbols represent atoms that are isotopes? A) 14C and 14N B) 16O and 18O C) 131I and 131I D) 222Rn and 222Ra Element Q has two isotopes. If 77% of the element has an isotopic mass of 83.7 amu and 23% of the element has an isotopic mass of 89.3 amu, what is the average atomic mass of the element? 1st % (1st mass #) + 2nd % (2nd mass #) + … 100 = 77(83.7) + 23(89.3) 100 = 84.988 amu (answer should be in between the two given masses, but closer to the one with the larger percent abundance) Definitions: Principal Energy Level: the main energy level or shell of an atom. Orbital: the region where an electron is most likely to be found. _____19. I can define principal energy level, orbital, ground state, excited state, electron configuration, and bright line spectrum. Ground State: electrons are in the lowest energy levels possible (closest to the nucleus); the Periodic Table shows ground state configurations. Excited State: a temporary state that exists when an electron jumps to a higher energy level (farther from the nucleus). Electron Configuration: shows the number of electrons in each shell. Bright Line Spectrum: narrow lines of color given off when an electron in the excited state releases energy and returns to the ground state. The 1st shell holds a maximum of 2 electrons. _____20. I can state the maximum number of electrons that will fit into each of the first four shells. The 2nd shell holds a maximum of 8 electrons. The 3rd shell holds a maximum of 18 electrons. The 4th shell holds a maximum of 32 electrons. _____21. I can define valence electrons. Definition: Valence Electrons: electrons in the last shell; it is the last number written in the electron configuration on the Periodic Table. _____22. I can define a stable octet. Definition: Stable Octet: having a full outer shell of 8 valence electrons; it makes the atom stable and unreactive. How many valence electrons does rubidium have in the ground state? 1 _____23. I can locate and interpret an element’s electron configuration on the Periodic Table. How many principal energy levels contain electrons in an atom of iodine in the ground state? 5 How many electrons does an aluminum atom have in its 2nd shell? 8 _____24. I can state the relationship between distance from the nucleus and energy of an electron. _____25. I can state the relationship between the number of the principal energy level and the distance to the atom’s nucleus. As the distance between the nucleus and the electron increases, the energy _____26. I can identify an electron configuration that shows an atom in the excited state. Which electron configuration represents an atom of potassium in the excited state? _____27. I can explain, in terms of subatomic particles and energy states, how a bright line spectrum is created. A bright-line spectrum is created when an electron in the excited state (a higher energy state) releases energy in the form of color as it returns to the ground state (a lower energy state). of the electron increases. As the number of the principal energy level increases, the distance to the nucleus increases. A) 2-8-7-1 B) 2-8-8-1 C) 2-8-7-2 D) 2-8-8-2 _____28. I can identify the elements shown in a bright line spectrum. Which element(s) is/are present in the mixture? D and E Draw the Lewis electron dot diagram for the following atoms: _____29. I can draw Lewis electron dot diagrams for a given element. Unit 3: Nuclear Chemistry Place a checkmark next to each item that you can do. If a sample problem is given, complete it as evidence. _____1. I can explain what makes a nucleus stable or unstable. The stability of the nucleus is dependent on the neutron to proton ratio. Type Symbol Notation alpha _____2. I can compare types of radiation. Charge Mass Penetrating Power +2 4 amu Low beta β- -1 0 amu Moderate positron β+ +1 0 amu Moderate gamma γ 0 0 amu High The difference between natural transmutation and artificial transmutation is that in natural transmutation an unstable nucleus breaks apart on its own _____3. I can explain the difference between natural transmutation and artificial transmutation. becoming a different element and in artificial transmutation a stable nucleus is made unstable by hitting it with a high energy particle (such as a proton, neutron, or gamma radiation). Once hit, the nucleus changes into a nucleus of a different element. Two synonyms for spontaneous decay are: natural transmutation _____4. I can state two synonyms for spontaneous decay. and natural decay. Which equation represents a natural decay? _____5. I can identify a natural decay reaction from a list of reactions. Which equation represents artificial transmutation? _____6. I can identify an artificial transmutation reaction from a list of reactions. Complete the following nuclear equations: _____7. I can show how mass number and electrical charge are conserved in any nuclear reaction. 30 15 P 92 36 _____8. I can define half-life. Kr Definition: Half-life: the amount of time it takes for half of the mass of a radioactive sample to decay. Based on Reference Table N, what fraction of a radioactive sample of Au198 will remain unchanged after 10.78 days? 1 half-life = 2.695 days, so 10.78 days is 4 half-lives = 1/16 of the sample (1 HL = 1/2, 2 HL = 1/4, 3 HL = 1/8, 4 HL = 1/16, 5 HL = 1/32) _____9. Given the length of the half-life and the amount of time that has passed, I can determine the amount of radioactive sample that will remain or the original amount. What was the original mass of a radioactive sample of K-37 if the sample decayed to 25.0 g after 4.92 seconds? H A T 0 400 g 0 1 200 g 1.23 s 2 100 g 2.46 s 3 50 g 3.69 s 4 25 g 4.92 s (half-life on Table N, add 1 half-life down T column, multiply by 2 to get to the original amount in between the two zeros) _____10. Given the length of the half-life and the amount of radioactive sample remaining, I can determine the amount of time that has passed. A 100.0 g sample of Co-60 decays until only 12.5 g of it remains. Given that the half-life of Co-60 is 5.271 years, how long did the decay take? H A T 0 100 g 0 1 50 g 5.271 y 2 25 g 10.542 y 3 12.5 g 15.813 y (divide by 2 to move down the A column, then add 1 half-life) _____11. Given the amount of time that has passed and the amount of radioactive sample remaining, I can determine the length of the half-life. What is the half-life of a radioisotope if 25.0 g of an original 200.0 g sample remains unchanged after 11.46 days? H A T 0 200 g 0 1 100 g 3.82 d 2 50 g 7.64 d 3 25 g 11.46 d 11.46 d / 3 (divide the last time by the number of half-lives that passed) Compared to K-37, the isotope K-42 has _____12. Using Table N, I can determine the length of half-life and/or decay mode for a specific radioactive isotope. A) shorter half-life and the same decay mode B) shorter half-life and a different decay mode C) longer half-life and the same decay mode D) longer half-life and a different decay mode _____13. I can explain why all nuclear reactions release A LOT more energy than chemical reactions do. Nuclear reactions release A LOT more energy than chemical reactions do because some of the mass is converted into energy. Definitions: Fission: when an atom is hit with a neutron, producing lighter nuclei, more neutrons, and energy (BIG SMALL). Fission equations usually contain uranium and neutrons are on both sides of the arrow. The neutrons produced allow a chain reaction to occur. _____14. I can define fission and fusion. Fusion: when two light nuclei combine to form a heavier nucleus (SMALL BIG). Fusion reactions have two hydrogen atoms, two helium atoms, or one hydrogen atom and one helium atom on the left side of the arrow. Which equation represents fission? _____15. I can identify a fission reaction from a list of reactions. Which equation represents fusion? _____16. I can identify a fusion reaction from a list of reactions. _____17. I can state the conditions of temperature and pressure that are needed for a fusion reaction to happen. The temperature and pressure conditions needed for fusion to happen are: high temperature and high pressure. Which of the following equations represent NUCLEAR reactions? _____18. Given a list of reactions, I can differentiate a nuclear reaction from a chemical reaction. Five beneficial uses for radioactive isotopes are: A. Tracing chemical and biological processes _____19. I can state 5 beneficial uses for radioactive isotopes. B. Detection and treatment of disease (must have short half-lives) C. Food storage D. Radioactive dating E. Nuclear power Tc-99 and Co-60 are used for treating cancer. _____20. I can state the scientific use of specific radioactive isotopes. I-131 is used for treating thyroid disorders. C-14 is used for dating once living organisms (wood, bone, animal skin). U-238 is used for dating geological formations. Three risks associated with radioactivity and radioactive isotopes are: A. Biological exposure and damage _____21. I can state three risks associated with radioactivity and radioactive isotopes. B. Long term storage and disposal (waste stays radioactive for long periods of time because they have very long half-lives) C. Nuclear accidents and pollution Unit 4: Periodic Table Place a checkmark next to each item that you can do. If a sample problem is given, complete it as evidence. _____1. I can state how the elements on the Periodic Table are arranged. The elements on the Periodic Table are arranged by increasing atomic number. Group 1 is called the alkali metals. Group 2 is called the alkaline earth metals. _____2. I can state the group names for elements in groups 1, 2, 3-12, 17, and 18. Groups 3-12 are called the transition metals or transition elements. Group 17 is called the halogens. Group 18 is called the noble gases. _____3. I can explain why elements in the same group have similar chemical properties. Elements in the same group have similar chemical properties because they have the same number of valence electrons (the last number in their electron configuration is the same). _____4. I can explain what elements in the same period have in common. Elements in the same period have the same number of occupied principal energy levels (shells) (their electron configurations consist of the same amount of numbers). Definitions: Electronegativity: the strength of attraction between the nucleus and a bonded electron; the higher the EN, the stronger the attraction (Table S). _____5. I can define electronegativity, ionization energy, atomic radius, and ionic radius. Ionization Energy: the amount of energy needed to remove the outermost electron (Table S). Atomic Radius: the distance from the nucleus to the outermost electron (the size of the atom) (Table S). Ionic Radius: the size of an ion once the electron(s) have been lost or gained. Classify each of the following elements as metals (M), nonmetals (NM), or metalloids (MTLD) (have properties of both metals and nonmetals). _____6. I can classify elements as metals, nonmetals, or metalloids based on their placement on the Periodic Table. MTLD B M K M Li NM C NM Ar MTLD Sb NM H M Fe M Au NM S MTLD Si M Fr NM He NM Rn M Al MTLD As M Bi NM I NM F MTLD Ge _____7. I can state the names/symbols for the two elements on the Periodic Table that are liquids at STP. The two elements that are liquids at STP are: Mercury (Hg) and Bromine (Br). The 11 elements that are gases at STP are: _____8. I can state the names/symbols of the 11 elements that are gases at STP. Helium (He), Neon (Ne), Argon (Ar), Krypton (Kr), Xenon (Xe), Radon (Rn), Oxygen (O), Nitrogen (N), Hydrogen (H), Fluorine (F), and Chlorine (Cl). The properties of metals are: Located to the left of the “staircase,” magnetic, malleable, ductile, have _____9. I can list properties of metals. luster, solids (except Mercury), good conductors of heat and electricity, high melting and boiling points, tend to lose electrons and form positive ions because they have large radii, low EN values, and low IE values. The properties of nonmetals are: Located to the right of the “staircase,” exist in all 3 phases (Group 17 contains all 3 phases), brittle, dull, poor conductors of heat and _____10. I can list properties of nonmetals. electricity (insulators), low melting and boiling points, tend to gain electrons and form negative ions because they have small radii, high EN values, and high IE values. As one reads down a group from top to bottom, the activity/reactivity of _____11. I can state the trend for activity/reactivity for METALS as one reads down a group. _____12. I can state the trend for activity/reactivity for NONMETALS as one reads down a group. METALS increases. The most reactive metal is Francium (Fr). As one reads down a group from top to bottom, the activity/reactivity of NONMETALS decreases. The most reactive nonmetal is Fluorine (F). Metals lose electrons and become positive ions that are smaller _____13. I can explain how losing or gaining electrons affects the radius of an ion. than the original atom because the number of shells decreases. Nonmetals gain electrons and become negative ions that are larger than the original atom because of more electron repulsion (more e -). _____14. I can explain why the elements in Group 18 don’t usually react with other elements. Elements in Group 18 don’t usually react with other elements because they have a stable octet of 8 valence electrons. Definition: Allotrope: forms of the same element in the same state that have different structures and different properties. _____15. I can define allotrope and give examples. Ex: Oxygen (O2 (g)) and Ozone (O 3 (g)) Ex: Carbon – diamond and graphite As one reads down a group from top to bottom, electronegativity decreases because the number of shells increases (shielding effect – inner _____16. I can state the periodic trend for electronegativity and explain why it occurs. shells block/weaken attraction between nucleus and valence electrons). As one reads across a period from left to right, electronegativity increases because the number of protons (nuclear charge) increases (more protons pull electrons in closer and strengthen the attraction). As one reads down a group from top to bottom, ionization energy decreases because the number of shells increases (shielding effect – inner shells block/weaken attraction between nucleus and valence electrons so _____17. I can state the periodic trend for ionization energy and explain why it occurs. it is easier for them to be removed). As one reads across a period from left to right, , first ionization energy increases because the number of protons (nuclear charge) increases (more protons pull electrons in closer and strengthen the attraction so it is harder for them to be removed). As one reads down a group from top to bottom, atomic radius increases because the number of shells increases (more shells take up _____18. I can state the periodic trend for atomic radius and explain why it occurs. more space). As one reads across a period from left to right, atomic radius decreases because the number of protons (nuclear charge) increases (more protons pull electrons in closer, making the atom smaller). Unit 5: Bonding Place a checkmark next to each item that you can do. If a sample problem is given, complete it as evidence. Definitions: Intramolecular Forces: forces of attraction within the same molecule. Examples: ionic bonds, covalent bonds, metallic bonds _____1. I can define intramolecular forces and intermolecular forces and give examples of each. Intermolecular Forces: forces of attraction between different molecules. Examples: Van der Waals (dipole-dipole forces, dispersion forces), hydrogen bonding, molecule-ion attraction. BARF stands for Break Absorb Release Form This means that when a bond is FORMED, energy is released and when a bond is BROKEN, energy is absorbed. _____ 2. I can explain and apply the meaning of BARF as it applies to chemical bonding. Given the balanced equation: N + N N2 Which statement describes the process represented by this equation? A) A bond is formed as energy is absorbed. B) A bond is formed as energy is released. C) A bond is broken as energy is absorbed. D) A bond is broken as energy is released. _____3. I can state the number of valence electrons that an atom attains to be most stable. _____4. I can state the three types of chemical bonds. Atoms are most stable when they have 8 valence electrons. The three types of chemical bonds are ionic, covalent, and metallic. Definitions: Ionic Bond: form between a metal and a nonmetal. Electrons are transferred from the metal to the nonmetal. The metal becomes a positive ion and the nonmetal becomes a negative ion and the two oppositely charged ions attract. _____5. I can define ionic bond, covalent bond, and metallic bond in terms of the types of elements (metals, nonmetals) from which they are formed. Covalent Bond: form between two or more nonmetals. Electrons are shared between the atoms. Metallic Bond: form between the same metal. The metals lose their electrons and the positive ions that form are attracted to the “sea of mobile electrons.” In an ionic bond, the valence electrons of the metals are transferred to the nonmetal so that each atom attains a stable octet (like noble gases). In these _____6. I can define ionic, covalent, and metallic bonds based on what happens to the valence electrons. compounds, the electronegativity difference is large (usually greater than 1.7). In a covalent bond, the valence electrons of the two nonmetals are shared so that each atom attains a stable octet. In these compounds, the electronegativity difference is small (usually less than 1.7). In a metallic bond, positive ions are immersed in a sea of mobile electrons. _____7. I can determine the compound that has the greatest ionic character. Which of the following compounds has the greatest ionic character? A) LiCl B) CaCl2 C) BCl3 D) RbCl Draw Lewis dot diagrams for the following ionic compounds: NaCl _____8. I can draw a Lewis dot diagram to represent an ionic compound. CaCl2 Physical properties of ionic substances are: they have hard, crystalline _____9. I can state the properties of ionic substances. structures as a solid (crystal lattice – geometric pattern), they have high melting and boiling points, they conduct electricity as a liquid and in solution (aq) (mobile charges), they are soluble (dissolve in water). A solid substance was tested in the lab. The results are shown below. *dissolves in water _____10. I can identify a substance as ionic based on its properties. *is an electrolyte * has a high melting point Based on these results, the solid substance could be A) Hg B) AuCl C) CH4 D) C12H22O11 Based on bond type, which compound has the highest melting point? A) CH4 B) C12H22O11 C) NaCl D) C5H12 In a single covalent bond, 2 electrons or 1 pair is shared. _____ 11. I can state the number of electrons that are shared in single and multiple covalent bonds. In a double covalent bond, 4 electrons or 2 pairs are shared. In a triple covalent bond, 6 electrons or 3 pairs are shared. Draw Lewis dot diagrams for the following molecular substances and identify their shapes. (Use the CNOF table) H2O - bent _____12. I can draw a Lewis dot diagram to represent a molecular (covalently bonded) substance and identify its shape as linear, bent, trigonal pyramidal, or tetrahedral. I2 - linear NH3 – trigonal pyramidal CO2 - linear CH4 - tetrahedral CHCl3 – tetrahedral Physical properties of molecular substances are: they can exist in all 3 _____13. I can state the properties of molecular substances. phases, they are usually not soluble in water, solids are soft, they are poor conductors in any phase (no mobile charges), they have low melting and boiling points. _____14. I can identify a substance as “molecular” based on its properties. _____15. I can state the type of bonding that occurs in compounds containing polyatomic ions. _____16. Given the chemical formula for a compound, I can determine the type(s) of bonding in the compound. Compounds containing polyatomic ions have both ionic and covalent bonding. State the type(s) of bonding in the following compounds: NaCl - ionic CO - covalent Hg - metallic Na3PO4 - ionic and covalent Explain, in terms of valence electrons, why the bonding in methane (CH4) is similar to the bonding in water (H2O). _____17. In terms of valence electrons, I can find similarities and differences between the bonding in several substances. In both molecules, the valence electrons are shared between the two atoms to form a covalent bond. Explain, in terms of valence electrons, why the bonding in HCl is different than the bonding in NaCl. In HCl the valence electrons are shared to form a covalent bond, while in NaCl the valence electrons are transferred to form an ionic bond. Polar covalent bonds are formed when different nonmetals share electrons _____18. I can explain the difference between a polar covalent bond and a nonpolar covalent bond. unequally. The electronegativity difference is between 0 and 1.7. Nonpolar covalent bonds are formed when the same nonmetals share electrons equally. The electronegativity difference is zero. _____19. I can explain how to determine the degree of polarity of a covalent bond. _____20. I can explain why one covalent bond is more or less polar than another covalent bond. The degree of polarity of a covalent bond is determined by the electronegativity difference between the elements. Explain, in terms of electronegativity difference, why the bond between carbon and oxygen in a carbon dioxide molecule is less polar than the bond between hydrogen and oxygen in a water molecule. The difference in electronegativity between carbon and oxygen is less than the difference in electronegativity between hydrogen and oxygen. Draw a water molecule and include slight charges. Explain why you assigned each atom its charge. _____21. I can determine which end of a polar covalent bond is slightly negative and which end is slightly positive and explain why. Oxygen is slightly negative because it has a higher electronegativity and hydrogen is slightly positive because it has a lower electronegativity. _____22. I can define symmetrical and asymmetrical. Definitions: Symmetrical: a molecule in which all slight charges cancel (looks the same when the molecule is turned upside down). There is an equal distribution of charge in a symmetrical molecule. Asymmetrical: a molecule in which all slight charges do not cancel (looks different when the molecule is turned upside down). There is an unequal distribution of charge in an asymmetrical molecule. SNAP means symmetrical nonpolar asymmetrical polar. Why is a molecule of CH4 nonpolar even though the bonds between the carbon and hydrogen are polar? _____23. I can explain and apply the meaning of SNAP as it relates to determining molecule polarity. A) The shape of the CH4 molecule is symmetrical. B) The shape of the CH4 molecule is asymmetrical. C) The CH4 molecule has an excess of electrons. D) The CH4 molecule has a deficiency of electrons. Explain, in terms of charge distribution, why a molecule of water is polar. Water is asymmetrical so the slight charges that exist do not cancel out (the molecule has asymmetrical distribution of charge). Determine which molecules are polar and which are nonpolar. Justify your answer. _____24. I can determine if a molecular is polar or nonpolar. H2O – polar, asymmetrical CO2 – nonpolar, symmetrical I2 – nonpolar, symmetrical (nonpolar bonds) CH4 – nonpolar, symmetrical NH3 – polar, asymmetrical CHCl3 – polar, asymmetrical Definitions: Dipole-Dipole: a type of van der Waal’s force that exists between polar molecules. The slight negative end of one polar molecule is attracted to the slight positive end of another polar molecule. _____25. I can define the different types of intermolecular forces. Dispersion: a type of van der Waal’s force that exists between nonpolar molecules. The attraction comes from the temporary slight charges that exist when the electrons move around the nucleus. Wherever the electron is, it brings a slight negative charge with it. Hydrogen Bonding: occurs between the hydrogen atom in one molecule and a nitrogen, oxygen, or fluorine atom in another molecule (H-NOF). Hydrogen bonds are very strong and explain the unusually high boiling point of water. Molecule-Ion Attraction: exists between ions of an ionic compound and the molecules of a polar liquid. This attraction is present in all solutions (aq) and explains how a substance dissolves. _____26. I can state the relationship between polarity and intermolecular force strength. _____27. I can state the relationship between size of the molecule and intermolecular force strength. _____28. I can state the relationship between the strength of the intermolecular forces and vapor pressure. _____29. Given the physical state of some substances, I can compare the relative strength of the intermolecular forces. _____30. Given the boiling points (or freezing points) of some substances, I can compare the relative strength of the intermolecular forces. As the polarity of the molecule increases, the strength of the forces increases. As the size of the molecule increases, the strength of the forces increases. As the strength of the forces increases, vapor pressure decreases. At STP, iodine (I2) is a crystal and fluorine (F2) is a gas. Compare the strength of the intermolecular forces in a sample of I2 at STP to the strength of the intermolecular forces in a sample of F2 at STP. Since I2 is a solid, it must have stronger intermolecular forces of attraction compared to fluorine, which is a gas. The stronger the forces of attraction, the closer the particles are arranged. At STP, CF4 boils at -127.8˚C and NH3 boils at -33.3˚C. Which substance has stronger intermolecular forces? Justify your answer. NH3 has stronger intermolecular forces of attraction because it has a higher boiling point. The stronger the forces of attraction, the higher the boiling point is. Which compound has hydrogen bonding between its molecules? A) CH4 _____31. I can answer questions about hydrogen bonding. B) CaH2 C) KNO3 D) H2O The relatively high boiling point of water is due to water having A) hydrogen bonding C) nonpolar covalent bonding B) metallic bonding D) strong ionic bonding In which system do molecule-ion attractions exist? A) NaCl (aq) B) NaCl (s) C) C6H12O6 (aq) D) C6H12O6 (s) Which diagram best illustrates the ion-molecule attractions that occur when the ions of NaCl (s) are added to water? _____32. I can answer questions about molecule-ion attraction. Unit 6: Chemical Reactions Place a checkmark next to each item that you can do. If a sample problem is given, complete it as evidence. Write the name for the following compounds: CrO chromium (II) oxide _____1. Given the chemical formula, I can write the name for binary compounds. MgI2 magnesium iodide Ni2O3 nickel (III) oxide FeCl2 iron (II) chloride SnS tin (II) sulfide Write the chemical formula for the following compounds: sodium bromide NaBr _____2. Given the name, I can write the chemical formula for binary compounds. lithium iodide LiI iron (III) fluoride FeF3 vanadium (V) oxide V2O5 Write the name for the following compounds: _____3. Given the chemical formula, I can write the name for ternary compounds. CuSO4 copper (II) sulfate (NH4)3PO4 ammonium phosphate Fe3(PO4)2 iron (II) phosphate Write the chemical formula for the following compounds: _____4. Given the name, I can write the chemical formula for ternary compounds. calcium oxalate CaC2O4 nickel (II) thiosulfate NiS2O3 Write the name for the following compounds: _____5. Given the chemical formula, I can write the name for covalent molecules. N2O dinitrogen monoxide SF6 sulfur hexafluoride N2O3 dinitrogen trioxide Write the chemical formula for the following compounds: _____6. Given the name, I can write the chemical formula for covalent molecules. diphosphorus pentasulfide P2S5 carbon monoxide CO dinitrogen tetroxide N2O4 How many atoms are in N2? 2 _____7.Given the chemical symbol/formula, I can determine how many atoms are present. What is the total number of atoms in Pb(C2H3O2)2? 15 How many atoms of carbon are in Pb(C2H3O2)2? 4 Using the symbols shown below, complete the equation below to illustrate conservation of mass. _____8. I can use particle diagrams to show conservation of mass in a chemical equation. = Al = Br 2 Al + 3 Br2 2 AlBr3 _____9. I can balance a chemical equation showing conservation of mass using the lowest whole number coefficients. Balance the following equation using the lowest whole number coefficients. _____10. Given a partially balanced equation, I can predict the missing reactant or product. Use the law of conservation of mass to predict the missing product. _____11. Given a list of chemical reactions, I can classify them as being a synthesis reaction, decomposition reaction, single replacement reaction, or double replacement reaction. _____12. From a list of given list of elements, I can determine which element is most active. Al2(SO4)3 + 3 Ca(OH)2 2 Al(OH)3 + 3 CaSO4 2 NH4Cl + CaO 2 NH3 + H2O + CaCl2 Classify the following reactions as synthesis, decomposition, single replacement, or double replacement. SR S D DR Which of the following elements is most likely to react? (Table J) A) Cu B) Al D) Mg C) Li Which reaction occurs spontaneously? _____13. I can predict if a single replacement reaction will occur. _____14. I can predict the products in a double replacement reaction. A) Cl2 + 2 NaBr Br2 + 2 NaCl C) I2 + 2 NaBr Br2 + 2 NaI B) Cl2 + 2 NaF F2 + 2 NaCl D) I2 + 2 NaF F2 + 2 NaI Complete and balance the following reaction: MgCl2 + 2 AgNO3 Mg(NO3)2 + 2 AgCl Unit 7: Moles Place a checkmark next to each item that you can do. If a sample problem is given, complete it as evidence. _____1. I can define gram formula mass (molar mass). The gram formula mass of a substance is the mass of 1 mole. _____2. I can define a mole as it pertains to chemistry. Definition: Mole: a unit of measurement for a substance; 1 mole = 6.02 x 1023 particles, mass is equal to the gram formula mass. 1 mole of a substance will occupy 22.4 L. What is the gfm for Ni2O3? Ni: 58.7 x 2 = 117.4 O: 16.0 x 3 = 48.0 gfm = 165.4 g _____3. I can determine the gram-formula mass for any compound or hydrate. What is the gfm for Pb(C2H3O2)2? Pb: 207.2 x 1 = 207.2 C: 12.0 x 4 = 48.0 H: 1.0 x 6 = 6.0 O: 16.0 x 4 = 64.0 gfm = 325.2 g What is the gfm for CuSO4 • 5 H2O? Cu: 63.5 x 1 = 63.5 S: 32.0 x 1 = 32.0 O: 16.0 x 4 = 64.0 H2O: 18.0 x 5 = 90.0 gfm = 249.5 g _____4. I can determine the percent composition of an element in a compound. What is the percent by mass of Mg in Mg(NO3)2? Mg: 24.3 x 1 = 24.3 % comp = part x 100 N: 14.0 x 2 = 28.0 whole O: 16.0 x 6 = 96.0 gfm = 148.3 g = 24.3 x 100 = 16.4 % 148.3 _____5. I can define hydrate. Definition: Hydrate: crystals with water inside; an anhydrous substance does not contain water and is formed when a hydrate is heated. _____6. I can determine the percent composition of water in a hydrate. What is the percent by mass of H2O in Na2SO4 • 10 H2O? Na: 23.0 x 2 = 46.0 % comp = part x 100 S: 32.0 x 1 = 32.0 whole O: 16.0 x 4 = 64.0 H2O: 18.0 x 10 = 180.0 = 180.0 x 100 = 55.9 % gfm = 322.0 g 322.0 A student heated a 9.1 g sample of a hydrate to a constant mass of 5.41 g. What percent by mass of water did the hydrate contain? Mass of H2O = hydrate – anhydrous = 9.1 – 5.41 = 3.69 % comp = part x 100 whole = 3.69 x 100 = 40.5 % 9.1 _____7. I can define empirical formula and molecular formula. _____8. Given the molecular formula, I can determine the empirical formula of a compound. _____9. Given the percent composition, I can determine the empirical formula of a compound. _____10. Given the empirical formula and the molar mass, I can determine the molecular formula of a compound. Definitions: Empirical Formula: the simplest ratio in which atoms combine; the subscripts are reduced. Molecular Formula: the actual ratios of atoms that combine; the subscripts are not reduced and are a multiple of the empirical. Which pair consists of a molecular formula and its corresponding empirical formula? A) C2H2 and CH3CH3 C) P4O10 and P2O5 B) C6H6 and C2H2 D) SO2 and SO3 A compound consists of 25.9% nitrogen and 74.1% oxygen by mass. What is the empirical formula of the compound? N = 25.9 g O = 74.1 g 14.0 16.0 = 1.85 = 4.63 N2O5 1.85 1.85 = 1x 2 = 2.5 x 2 What is the molecular formula of a compound that has the empirical formula of CH2O and a molar mass of 180 g/mol? C: 12.0 x 1 = 12.0 H: 1.0 x 2 = 2.0 C6H12O6 O: 16.0 x 1 = 16.0 gfm = 30.0 g molecular mass = 180 = 6 empirical mass 30 _____11. I can find the number of moles of a substance if I am given the mass and formula for the substance. How many moles of NaCl are equivalent to 94.3 g? Na: 23.0 x 1 = 23.0 moles = mass x = 94.3 = 1.6 moles Cl: 35.5 x 1 = 35.5 gfm 58.5 gfm = 58.5 g _____12. I can find the mass of a substance if I am given the moles and formula for the substance. How many grams of LiBr would 3.5 moles of LiBr be? Li: 6.9 x 1 = 6.9 moles = mass 3.5 = x Br: 79.9 x 1 = 79.9 gfm 86.8 gfm = 86.8 g How many moles of O2 contain 3.01 x 1023 molecules? _____13. I can convert between moles and numbers of particles using Avogadro’s number. 1 mole = x 6.02 x 1023 3.01 x 1023 x = 0.5 moles How many molecules are in 0.25 moles of H2? 1 mole = 6.02 x 1023 0.25 x x = 1.5 x 1023 molecules x = 303.8 g How many liters does 4.6 moles of N2 occupy? 1 mole 22.4 L _____14. I can convert between moles and volume. _____16. I can define stoichiometry. 4.6 x x = 103.04 L How many moles of CO2 are present in an 11.2 L sample? 1 mole 22.4 L _____15. I can answer questions using Avogadro’s Hypothesis. = = x 11.2 x = 0.5 moles Which gas sample at STP has the same total number of molecules as 2 L of CO2 (g) at STP? A) 5 L of CO2 (g) B) 2 L of Cl2 (g) C) 3 L of H2S (g) D) 6 L of He (g) Definition: Stoichiometry: a proportional relationship between two or more substances in a reaction. Given the following balanced equation, state the mole ratios between the following substances. C3H8 (g) + 5 O2 (g) 3 CO2 (g) + 4 H2O () _____17. Given a balanced equation, I can state the mole ratios between any of the reactants and/or products. The mole ratio between C3H8 and O2 is 1 to 5. The mole ratio between C3H8 and CO2 is 1 to 3. The mole ratio between C3H8 and H2O is 1 to 4. The mole ratio between CO2 and O2 is 3 to 5. The mole ratio between H2O and CO2 is 4 to 3. Using the equation from question #17, determine how many moles of O 2 are needed to completely react with 7.0 moles of C3H8. _____18. Given the number of moles of one of the reactants or products, I can determine the number of moles of another reactant or product that is needed to react. 5 O2 = 1 C3H8 x 7 x = 35 moles of O2 Using the equation from question #17, determine how many moles of CO2 are produced when 8.0 moles of H2O completely react. 3 CO2 4 H2O = x 8 x = 6 moles of CO2 Unit 8: Solutions Place a checkmark next to each item that you can do. If a sample problem is given, complete it as evidence. Definitions: Solute: the substance being dissolved; usually present in a smaller amount. Ex – salt in a salt water solution. _____1. I can define solute, solvent, solution, and solubility. Solvent: the substance that does the dissolving; usually present in a greater amount. Ex – water in a salt water solution. Solution: a homogeneous mixture of solute and solvent. Solubility: the amount of solute that will dissolve in a certain amount of solvent at a given temperature. “Like dissolves like” means the polarity of the solute and the polarity of the solvent must be the same in order for a substance to dissolve. Nonpolar solutes dissolve in nonpolar solvents. Polar or ionic solutes dissolve in polar solvents. _____2. I can explain and apply the expression “like dissolves like.” Explain, in terms of molecular polarity, why ammonia is more soluble than methane in water at 20˚C at standard pressure. Ammonia is a polar molecule and will dissolve in water which is also a polar molecule. Methane is a nonpolar molecule and will not dissolve in polar water. _____3. I can describe the trend in solubility for solids and liquids as the temperature changes. _____4. I can describe the trend in solubility for gases as the temperature changes. _____5. I can describe the trend in solubility for gases as the pressure changes. _____6. I can describe the conditions in which gases are most soluble. As the temperature increases, the solubility of solids and liquids increases. As the temperature increases, the solubility of a gas decreases. As the pressure increases, the solubility of a gas increases. Gases are most soluble in low temperatures and high pressures. Write “S” for soluble and “NS” for not soluble. _____7. I can use Table F to determine if a substance will be soluble in water. _____8. I can use Table G to determine how much solute to add at a given temperature to make a saturated solution. S potassium chlorate NS silver bromide S lithium carbonate NS calcium carbonate How many grams of KClO3 must be dissolved in 100 grams of water at 20˚C to make a saturated solution? 8 grams (must be on the KClO 3 line to be saturated) _____9. I can use Table G to determine if a solution is saturated, unsaturated, or supersaturated. If 20.0 g of NaNO3 are dissolved in 100.0 g of water at 25.0˚C, will the resulting solution be saturated, unsaturated, or supersaturated? Unsaturated (point below the line) Saturated solutions are on the line, supersaturated solutions are above. Definitions: Concentration: amount of dissolved solute in a given amount of solvent. _____10. I can define concentration, dilute, and concentrated. Dilute: a solution that contains a small amount of dissolved solute. Concentrated: a solution that contains a large amount of dissolved solute. What is the molarity of 3.5 moles of NaBr dissolved in 500 mL of water? Molarity = moles L _____11. I can calculate the concentration of a solution in molarity. x = 3.5 .5 x=7M What is the molarity of 1.5 L of a solution containing 52 g of LiF? moles = mass gfm Molarity = moles L x = 52 25.9 x=2 1.5 x = 2 moles x = 1.3 M Which solution is most concentrated? _____12. I can determine which solution is the most concentrated or the most dilute. _____13. I can calculate the concentration of a solution in ppm. A) 1 mole of solute dissolved in 1 liter of solution B) 2 moles of solute dissolved in 3 liters of solution C) 6 moles of solute dissolved in 4 liters of solution D) 4 moles of solute dissolved in 8 liters of solution What is the concentration, in ppm, of a 2,600 g solution containing 0.015 g of CO2? ppm = mass of solute x 1,000,000 ppm = 0.015 x 1,000,000 = 5.77 ppm mass of solution 2,600 _____14. I can explain how adding a solute to a pure solvent affects the freezing point of the solvent. When a solute is added to a solvent, the freezing point of the solvent ____15. I can explain how adding a solute to a pure solvent affects the boiling point of the solvent. When a solute is added to a solvent, the boiling point of the solvent ____16. I can determine which solution will have the lowest freezing point/highest boiling point. Which solution has the highest boiling point? decreases. The more solute that is added, the lower the freezing point gets. increases. The more solute that is added, the higher the boiling point gets. A) 1.0 M KNO3 B) 2.0 M KNO3 C) 1.0 M Ca(NO3)2 D) 2.0 M Ca(NO3)2 Unit 9: Acids & Bases Place a checkmark next to each item that you can do. If a sample problem is given, complete it as evidence. _____1. I can use two different systems to define acids and bases. _____2. I can determine which substance is an acid or base according to the alternative theory given a reaction. Arrhenius Theory Alternative Theory (Bronsted-Lowry) Acid Produces H+ ions (H3O+) Donates H+ ions Base Produces OH- ions Accepts H+ ions Given the reaction: HI + H2O H3O+ + I-, identify the acid and base. Acid: HI Base: H2O List the chemical names of three common acids and three common bases. _____3. I can give examples of the chemical names of common acids and bases. Acids (Table K) Bases (Table L) Hydrochloric Acid Sodium Hydroxide Nitric Acid Potassium Hydroxide Sulfuric Acid Calcium Hydroxide List the chemical formulas of three common acids and three common bases. _____4. I can give examples of the chemical formulas of common acids and bases. _____5. I can list properties of common acids and bases. Acids (Table K) Bases (Table L) HCl NaOH HNO3 KOH H2SO4 Ca(OH)2 Acids Taste sour, electrolytes, react with bases to form water and salt in a neutralization reaction, react with active metals (higher than H2 on Table J) to produce hydrogen gas and salt, cause indicators to change color Bases Taste bitter, feel slippery, electrolytes, react with acids to form water and salt in a neutralization reaction, do not react with metals, cause indicators to change color Definitions: Electrolyte: any substance that conducts electricity when dissolved in water; produces mobile charges. _____6. I can define electrolyte, nonelectrolyte, hydronium ion, and hydroxide ion. Nonelectrolyte: any substance that does not conduct electricity when dissolved in water; does not produce mobile charges. Hydronium Ion: H3O+, also known as a hydrogen ion; is produced by acids and is present in larger amounts in acidic solutions. Hydroxide Ion: OH-; is produced by bases and is present in larger amounts in basic solutions. _____7. I can state another name for the hydronium ion. _____8. I can state the two types of substances that are nonelectrolytes. _____9. I can state the three types of substances that are electrolytes. _____10. I can identify nonelectrolytes. _____11. I can identify electrolytes. The hydronium ion is also known as the hydrogen ion (H+). Sugars and alcohols are two classes of compounds that are nonelectrolytes. Acids, bases, and salts are three classes of compounds that are electrolytes. Circle all of the substances that are nonelectrolytes. A) C2H5OH B) C6H12O6 C) C12H22O11 D) CH3COOH E) LiOH Which substance, when dissolved in water, forms a solution that conducts an electric current? A) C2H5OH B) C6H12O6 C) C12H22O11 D) CH3COOH Predict the products in the following reactions, then balance. Zn + 2 HCl H2 + ZnCl2 _____12. I can complete reactions between metals and acids. 3 Mg + 2 H3PO4 3 H2 + Mg3(PO4)2 Cu + H2SO4 No Reaction _____13. I can define neutralization. Definition: Neutralization: a double replacement reaction that occurs between an acid and a base producing water and a salt. Which of the following equations is a neutralization reaction? _____14. I can identify a neutralization reaction from a list of reactions. A) 6 Na + B2O3 3 Na2O + 2 B B) Mg(OH)2 + 2 HBr MgBr2 + 2 H2O C) 2 H2 + O2 2 H2O D) 2 KClO3 2 KCl + 3 O2 Definitions: pH: the concentration of hydrogen ions in solution; measures how acidic or basic a solution is. _____15. I can define pH, pOH, and [ ]. pOH: the concentration of hydroxide ions in solution. [ _____16. I can describe the relationship between pH and pOH and [H+] and [OH-]. ]: concentration pH + pOH = 14 [H+] and [OH-] are indirectly related. If the pH of a solution is 4.5, the solution is acidic. _____17. Based on pH, I can determine if a solution is acidic, basic, or neutral. If the pH of a solution is 7.0, the solution is neutral. If the pH of a solution is 11, the solution is basic. If the pH of a solution is 5.7, the solution is acidic. _____18. I can identify the pH values of strong acids and strong bases. Strong acids have low pH values and strong bases have high pH values. In an acidic solution the [H+] is greater than the [OH-]. _____19. I can describe acidic, basic, or neutral solutions in terms of ion concentrations. In a basic solution the [H+] is less than the [OH-]. In a neutral solution the [H+] is equal to the [OH-]. As the H+ concentration decreases, the pH increases. _____20. I can state the relationship between H+ concentration and pH. As the H+ concentration increases, the pH decreases. A greater number of hydrogen ions results in a lower pH, indicating a strong acid. If the [H3O+] is 1 x 10-8, the pH of the solution will be 8. If the [H3O+] is 1 x 10-1, the pH of the solution will be 1. _____21. Given the hydronium ion concentration, I can determine the pH. If the [H3O+] is 1 x 10-14, the pH of the solution will be 14. If the [H3O+] is 1 x 10-7, the pH of the solution will be 7. If the pH is 10, the [H+] is 1 x 10-10 M. _____22. Given the pH, I can determine the [H+] or [OH-]. If the pH is 3, the [H+] is 1 x 10-3 M. If the pH is 6, the [OH-] is 1 x 10-8 M. If the pH is 9, the [OH-] is 1 x 10-5 M. If the H+ concentration is increased by a factor of 10, the pH will decrease by 1. _____23. I can determine the change in pH when the [H+] of a solution is changed. If the H+ concentration is increased by a factor of 100, the pH will decrease by 2. If the H+ concentration is decreased by a factor of 1,000, the pH will increase by 3. _____24. I can answer indicator questions using Table M. Phenolphthalein tests colorless and bromthymol blue tests blue in samples of fish tank water. What is the pH range for the water in this fish tank? 7.6 – 8.0 What color is thymol blue in a pH of 11? Blue _____25. I can state the name of the laboratory equipment that is used to carry out a titration. Which piece of laboratory equipment is used to carry out a titration? Burette _____26. I can state the purpose of titration. Why do scientists do titrations? To determine the concentration (molarity) of an unknown acid or base. In a titration, 36 mL of a 0.16 M NaOH solution exactly neutralizes 12 mL of an H2SO4 solution. What is the concentration of the H2SO4 solution? (H+)MAVA = MBVB(OH-) (2)(MA)(12) = (0.16)(36)(1) MA = 0.24 M A student uses a 1.4 M HBr solution in a titration to determine the concentration of a KOH solution. The data is below: _____27. I can solve for any variable in the titration equation. What is the molarity of the KOH solution? (H+)MAVA = MBVB(OH-) (1)(1.4)(15.4) = (MB)(22.1)(1) MB = 0.98 M (subtract the volumes to get the volume used – plug those numbers in) Unit 10: Kinetics & Equilibrium Place a checkmark next to each item that you can do. If a sample problem is given, complete it as evidence. _____1. I can define an effective collision according to the collision theory. _____2. I can determine the type of compound that has a faster reaction rate. Definition: Effective Collision: a collision between reactants that have enough energy and proper angles; results in a reaction. Ionic substances react faster than covalent molecules because their reactions involve less bond rearrangement. As the concentration increases, the reaction rate increases because there are _____3. I can state and apply the relationship between concentration and reaction rate in terms of the collision theory. more effective collisions between particles. At 20˚C, a reaction between powdered Zn(s) and hydrochloric acid will occur most quickly if the concentration of the HCl is A) 1.0 M B) 1.5 M C) 2.5 M D) 2.8 M As the surface area increases, the reaction rate increases because there are _____4. I can state and apply the more effective collisions between particles. relationship between surface area and reaction rate in terms of the At STP, which 4.0 g sample of Zn(s) will react most quickly with dilute collision theory. hydrochloric acid? A) lump B) bar C) powdered D) sheet metal As the temperature increases, the reaction rate increases because there are _____5. I can state and apply the relationship between temperature and reaction rate in terms of the collision theory. more effective collisions between particles. Given the reaction: 2 Mg (s) + O2 (g) 2 MgO (s) At which temperature would the reaction occur at the greatest rate? A) 0˚C _____6. I can use Table I to determine if a reaction is exothermic or endothermic. B) 15˚C C) 95˚C D) 273 K _____7. Based on the location of the energy term, I can determine if the reaction is exothermic or endothermic. _____8. I can determine the new heat of reaction if the number of moles in a reaction changes. Given the following balanced equation: I + I I2 + 146.3 kJ Is this reaction exothermic or endothermic? Justify your answer. Exothermic – energy is released (it is on the right of the arrow – product). Given the following balanced equation: 2 C + 3 H2 C2H6, what is the heat of reaction if 3 moles of carbon react? 2C = 3 -84 x x = -126 kJ Definitions: Potential Energy Diagram: a graph that shows the changes in potential energy over the course of a chemical reaction. _____9. I can define potential energy diagram, heat of reaction (∆H), activation energy, and catalyst. Heat of Reaction (∆H): the difference between the potential energy of the reactants and the potential energy of the products; Hproducts – Hreactants. Activation Energy: the minimum amount of energy needed for a chemical reaction to occur. Catalyst: a substance that makes a reaction occur faster by lowering the activation energy and providing an alternate pathway for the reaction to occur. Given the potential energy diagram below, determine the following: _____10. Given a potential energy diagram, I can determine the PE of reactants, PE of products, and PE of the activated complex, ∆H, and activation energy of forward and reverse reactions. PE of reactants = 40 kJ PE of products = 15 kJ PE of activated complex = 50 kJ ∆H= -25 kJ Activation Energy of Forward Reaction = 10 kJ Activation Energy of Reverse Reaction = 35 kJ _____11. Given a potential energy diagram for an uncatalyzed reaction, I can determine how the diagram will change when a catalyst is added. Draw a dotted line on the potential energy diagram below to indicate how it will change if a catalyst is added. List 3 values that will change and 3 values that will not change when a catalyst is added. Will Change: PE of activated complex, activation energy of forward reaction, activation energy of reverse reaction Will Not Change: PE of reactants, PE of products, heat of reaction (∆H) Give the potential energy diagram below, determine if the reaction is exothermic or endothermic. Justify your answer. Endothermic – energy of the products is greater which shows that energy was absorbed during the reaction. The difference in PE of the reactants and the PE of the products is positive. _____12. Given a potential energy diagram, I can determine if the reaction is exothermic or endothermic. In a system at equilibrium the rates of the forward and reverse reaction _____13. I can state two conditions that apply to all systems at equilibrium. must be equal and the concentrations of the reactants and products must be constant. Two types of physical equilibrium are phase equilibrium, in which phase _____14. I can describe examples of physical equilibrium. changes occur at the same rate, and solution equilibrium. In terms of saturation, a solution that is at equilibrium must be saturated. Definitions: Forward Reaction: the chemical reaction read from left to right. _____15. I can define forward reaction, reverse reaction, reversible reaction, and closed system. Reverse Reaction: the chemical reaction read from right to left. Reversible Reaction: a chemical reaction that can occur in both directions (forward and reverse). Closed System: a system in which reactants and products may not enter or leave. _____16. I can state LeChatelier’s Principle. LeChatelier’s Principle states when a reaction at equilibrium is subjected to a stress, it will shift (favor the forward or reverse reaction) to relieve the stress. Given the reaction at equilibrium: 2 SO2 (g) + O2 (g) _____17. Given a balanced equation at equilibrium, I can predict the direction of shift in the equilibrium when the concentration, temperature, or pressure is changed or if a catalyst is added. I can predict which substances will increase in concentration and which substances will decrease in concentration due to the shift. _____18. I can define entropy. _____19. I can rank the three phases of matter from least entropy to most entropy. 2 SO3 (g) + 392 kJ Predict the direction of shift in the equilibrium (right, left, none) when the following changes are made to the system. Then predict which substances from the above reaction will increase and decrease due to the shift. Increase [SO2] Right Increased Concentration of SO3 Increase [SO3] Left SO2, O2 SO3 Decrease [O2] Increase temperature Decrease temperature Increase pressure Left SO2, O2 SO3 Left SO2, O2 SO3 Right SO3 SO2, O2 Right SO3 SO2, O2 Decrease pressure Left SO2, O2 SO3 Add a catalyst None None None Change Direction of Shift Decreased Concentration of SO2, O2 Entropy is the amount of randomness or disorder. Least entropy Most entropy Solid < Liquid < Gas _____20. Given a balanced equation, I can determine if the reaction results in an overall increase or decrease in entropy. _____21. I can state the trends in nature for entropy and energy. Most systems in nature tend to undergo reactions that have a(n) increase in entropy and a(n) decrease in energy. Unit 11: Redox Place a checkmark next to each item that you can do. If a sample problem is given, complete it as evidence. _____1. I can assign oxidation numbers to any element or ion. Assign oxidation number to each of the elements below. O2 0 Li 0 Na+ +1 Assign oxidation numbers to each element in the compounds below. _____2. I can assign oxidation numbers to the elements in a compound. MnCl3: Mn +3 Cl -1 H2SO4: H +1 S +6 O -2 Assign oxidation numbers to each element in the polyatomic ions below. _____3. I can assign oxidation numbers to the elements in a polyatomic ion. _____4. I can state the Law of Conservation of Charge. PO4-3: P +5 O -2 ClO3-1: Cl +5 O -2 The Law of Conservation of Charge states in any reaction, the total charge of the reactants must be the same as the total charge of the products. Definitions: Oxidation: the loss of electrons; the oxidation number will increase. Reduction: the gain of electrons; the oxidation number will decrease. _____5. I can define oxidation, reduction, redox reaction, oxidizing agent, and reducing agent. Redox Reaction: a reaction in which there is a transfer of electrons; one substance will lose electrons and another substance will gain electrons. Both oxidation and reduction occur simultaneously. Oxidizing Agent: the substance that gets reduced. Reducing Agent: the substance that gets oxidized. _____6. I can identify a redox reaction from a list of chemical reactions. Which half-reaction equation represents the reduction of a potassium ion? A) K+ + 1 e- K0 C) K+ K0 + 1 eB) K0 + 1 e- K+ D) K0 K+ + 1 e- _____7. I can distinguish between an oxidation halfreaction and a reduction halfreaction. The two half-reactions that come from the following equation are: Li0 (s) + Ag+ (aq) Li+ (aq) + Ag0 (s) _____8. I can break a redox reaction into its two halfreactions. Oxidation Half-Reaction: Li0 (s) Li+ (aq) + 1eReduction Half-Reaction: Ag+ (aq) + 1e- Ag0 (s) Given the reaction: Mg0 (s) + Fe+3 (aq) Fe0 (s) + Mg+2 (aq) When the equation is correctly balanced using smallest whole numbers, the coefficient of Fe+3 will be A) 1 _____9. I can balance a redox reaction. B) 2 C) 3 D) 6 _____10. I can describe the conditions needed for a spontaneous redox reaction to occur. A spontaneous redox reaction will occur if the metal being oxidized is _____11. I can state the two types of electrochemical cells. The two types of electrochemical cells are: voltaic and electrolytic. higher on Table J, or the nonmetal reduced is higher on Table J. Voltaic Electrolytic One half-cell with the anode, one half-cell with the cathode, salt bridge One container with the anode and cathode, battery Oxidation Site Anode Anode Reduction Site Cathode Cathode Anode Charge Negative Positive Cathode Charge Positive Negative Mass at Anode Decreases Decreases Mass at Cathode Increases Increases Anode Cathode Anode Cathode Chemical Electrical Electrical Chemical Spontaneous Not Spontaneous Batteries Electroplating Components _____12. I can compare the two types of electrochemical cells. Direction of Electron Flow Energy Conversion Nature of Process Use _____13. I can state the purpose of the salt bridge in a voltaic cell. The purpose of the salt bridge is to allow for the migration of ions; it balances the charges and maintains neutrality in the solutions. _____14. I can answer questions about a voltaic cell. What is the anode? Zn0 What is the cathode? Cu0 Write the oxidation half-reaction: Zn 0 Zn+2 + 2 eWrite the reduction half-reaction: Cu+2 + 2 e- Cu 0 In what direction do electrons flow? Anode Cathode (Zn0 to Cu 0) Explain, in terms of atoms and ions, why the mass of the cathode increases during the operation of an electrochemical cell. _____15. I can explain, in terms of atoms and ions, the changes in mass that take place at the anode and cathode of an electrochemical cell. The ions from the solution are reduced and become atoms that become part of the cathode. Explain, in terms of atoms and ions, why the mass of the anode decreases during the operation of an electrochemical cell. The atoms from the anode are oxidized producing ions that move into the solution. The diagram below represents an electrochemical cell: _____16. I can answer questions about an electrolytic cell. What is the anode? Ag0 What is the cathode? Spoon Write the oxidation half-reaction: Ag0 Ag+ + 1 eWrite the reduction half-reaction: Ag+ + 1 e- Ag0 In what direction do electrons flow? Anode Cathode (Ag to spoon) What is the purpose of the battery? Provide energy for the reaction. Unit 12: Organic Chemistry Place a checkmark next to each item that you can do. If a sample problem is given, complete it as evidence. _____1. I can state the name and symbol of the element that is capable of forming rings, chains, and networks. The element that is capable of forming rings, chains, and networks is carbon. Its symbol is C. Definitions: Organic Compound: a compound containing a carbon-hydrogen covalent bond. _____2. I can define organic compound, saturated hydrocarbon, unsaturated hydrocarbon, isomers, and alkyl group. Saturated Hydrocarbon: a compound containing only carbon and hydrogen in which all carbon atoms are held together by single bonds. Unsaturated Hydrocarbon: a compound containing only carbon and hydrogen in which there is at least one double or triple bond (multiple bond) between the carbon atoms. Isomers: molecules with the same molecular formula, but they have different structures, properties, and names. Alkyl Group: a hydrocarbon branch off of the main chain. _____3. Given the formula, I can determine if a compound is a hydrocarbon. Draw the complete structural formula for CH3CH2CH2CH2CH3. _____4. I can expand a condensed structural formula to show the structural formula of an organic compound. _____5. Given the formula, I can determine which compound has the highest boiling point. _____6. Given the name or formula, I can use Table Q to determine which class of hydrocarbons a compound belongs. Draw the complete structural formula for CH3CHCHCH3. Which of the following hydrocarbons has the highest boiling point? A) C3H8 B) C4H10 C) CH4 D) C2H6 Determine the homologous series of hydrocarbons to which each of the following belongs: Decane: Alkane 2-Decene: Alkene C4H8: Alkene C7H12: Alkyne _____7. Given the name or formula, I can determine if the hydrocarbon is saturated or unsaturated. Determine if each of the following is a saturated or unsaturated hydrocarbon. Decane: saturated 2-Decene: unsaturated CH3CH2CH2CH3: saturated CH3CHCHCH3: unsaturated Determine the homologous series of hydrocarbons to which each of the following belongs: belongs to the alkane series. _____8. Given the structural formula, I can determine which homologous series a hydrocarbon belongs to. belongs to the alkyne series. belongs to the alkene series. Determine how many carbon atoms are in each of the following compounds: _____9. Given the name, I can use Table P to determine how many carbons atoms are in a compound. Decane: 10 Ethene: 2 3-Nonene: 9 1-Pentyne: 5 Propene: 3 Methane: 1 _____10. Given a list of compounds, I can determine which ones are isomers. Draw an isomer of the following: _____11. I can draw an isomer of a given compound. Name the following hydrocarbons: 2-pentene 3-heptyne nonane _____12. I can use Tables P & Q to name hydrocarbons. 2-hexene 2-methyl hexane 2,2,3-trimethyl butane Draw the following hydrocarbons: propyne _____13. I can use Tables P & Q to draw hydrocarbons. 3-ethyl hexane 3-hexene 2,3-dimethyl butane _____14. Given a structural formula, I can identify which class of organic compounds a substance belongs to. Name the following organic compounds: _____15. I can use Tables P & R to name simple compounds in any of the classes of organic compounds. butane butanone butanoic acid 2-butene 1-butanol 1-butanamine 2-butyne butanal butanamide 1-bromobutane methyl propanaote propyl methyl ether Draw the following organic compounds: 2,2-dichloropentane dimethyl ether _____16. I can use Tables P & R to draw simple compounds in any of the classes of organic compounds. 2-pentanone methyl butanoate pentanamide 2-butanol ethanal propanoic acid ethanamine ethyl propanoate Substitution (halogenation) is the replacement of a hydrogen in a saturated hydrocarbon with a halogen. It starts with an alkane and ends with two product(s). _____17. I can describe the 7 types of organic reactions. Addition is the addition of hydrogens or halogens at a double bond in an unsaturated hydrocarbon. It starts with an alkene and ends with one product(s). Esterification is when a(n) organic acid and a(n) alcohol react to form an ester and water. Saponification is the production of soap. Fermentation is the breakdown of sugar into alcohol and carbon dioxide. Combustion is when a hydrocarbon reacts with oxygen to produce water and carbon dioxide. Polymerization is the formation of a large molecule called a polymer from small molecules called monomers. When monomers are joined together by breaking double or triple bonds it is called addition polymerization. When monomers are joined by the removal of water it is called condensation polymerization. _____18. Given an equation, I can identify the type of organic reaction that is occurring. Complete the following reactions: _____19. I can complete an organic reaction by predicting the missing products.