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Science Curriculum Matrix Chemistry August 1, 2009 The Science Vertical Team has revised the Chemistry Curriculum Matrix for 2009-2010. In addition to the necessary correlation to the Virginia Science Standards of Learning, the Chemistry content is organized by both concepts and topics. We encourage you to utilize this document while planning for instruction. A more dynamic version of this matrix is available on our wiki site at http://acpsscience.pbworks.com/. We anticipate making additional updates to this document as the school year progresses. Please contact Tony Borash with your comments and suggestions at [email protected]. In addition to this document, we recommend that you review the Chemistry Curriculum Framework for additional clarification regarding the Chemistry Science SOL and the Chemistry Enhanced Scope and Sequence for unit and lesson planning resources. Thanks, The Science Vertical Team Scientific Investigation, Reasoning, and Logic The Science Vertical Team will work to develop a recommended integration for the science processes and skills generally found in standard CH.1 into the more content specific standards CH.2 through CH.5. In the meantime, Virginia Science SOL CH.1 and the associated essential skills and processes are presented below. Please incorporate these specific science processes and skills into your daily science instruction as much as possible and practical. CH.1 The student will investigate and understand that experiments in which variables are measured, analyzed, and evaluated, produce observations and verifiable data. a. b. c. d. e. f. g. h. i. designated laboratory techniques; safe use of chemicals and equipment; proper response to emergency situations; manipulation of multiple variables with repeated trials; accurate recording, organizing, and analysis of data through repeated trials; mathematical and procedural error analysis; mathematical manipulations (SI units, scientific notation, linear equations, graphing, ratio and proportion, significant digits, dimensional analysis); the use of appropriate technology including computers, graphing calculators, and probeware for gathering data and communicating results; and construction and defense of a scientific viewpoint (the nature of science). Essential Skills and Processes In order to meet this standard, it is expected that students should know and be able to: For CH.1 a,b, and c: • • • • • • Make the following measurements, using the specified equipment: - volume: graduated cylinder, pipette, volumetric flask, buret - mass: electronic or dial-a-gram - temperature: thermometer and/or temperature probe - pressure: barometer or pressure probe. Identify, locate, and know how to use laboratory safety equipment, including aprons, goggles, gloves, fire extinguishers, fire blanket, safety shower, eye wash, broken glass container, and fume hood. Demonstrate the following basic lab techniques: filtering, decanting, using chromatography, and lighting a gas burner. Identify the following basic lab equipment: beaker, flask, graduated cylinder, test tube, test tube rack, test tube holder, ring stand, wire gauze, clay triangle, crucible with lid, evaporation dish, watch glass, wash bottle, and dropping pipette. Understand Material Safety Data Sheet (MSDS) warnings, including handling chemicals, lethal dose (LD), hazards, disposal, and chemical spill cleanup. Demonstrate safe laboratory practices, procedures, and techniques. For CH.1 d and e: • • Design and perform experiments to test predictions. Identify variables. • • • • Predict outcome(s) when a variable is changed. Record data, using the significant digits of the measuring equipment. Demonstrate precision (reproducibility) in measurement. Recognize accuracy in terms of closeness to the true value of a measurement. For CH.1 f and g: • Discover and eliminate procedural errors. • Know most frequently used SI prefixes and their values (milli-, centi-, deci-, kilo-). • Demonstrate the use of scientific notation, using the correct number of significant digits with powers of ten notation for the decimal place. • Correctly utilize the following when graphing data. - dependent variable (vertical axis) - independent variable (horizontal axis) - scale and units of a graph - regression line (best fit curve). • Calculate mole ratios, percent composition, conversions, and relative atomic mass. • Use the rules for performing operations with significant digits. • Utilize dimensional analysis. • Use graphing calculators correctly. • Read a measurement from a graduated scale, stating measured digits plus the estimated digit. • Use data collected to calculate percent error. • Determine the mean of a set of measurements. For CH.1 h and i: • • • Use appropriate technology for data collection and analysis, including probeware interfaced to a graphing calculator and/or computer. Use probeware to gather data. Explain the emergence of modern theories based on historical development. For example, students should be able to explain the origin of the atomic theory beginning with the Greek atomists and continuing through the most modern Quantum models. Physical Science: Matter: Structure COURSE: Chemistry CONCEPT: SCALE: Properties ENDURING UNDERSTANDING: Properties characterize objects, organisms, and substances. Essential Understandings Assessment Samples – SOL/Blooms Students should understand: Knowledge/Comprehension Level • The periodic table is arranged in order of increasing atomic numbers. • The atomic number of an element is the same as the number of protons. • In a neutral atom, the number of electrons is the same as the number of protons. • All atoms of an element have the same number of protons. • The atomic mass for each element is the weighted average of that element’s naturally occurring isotopes. • Students use colored counters that represent protons, electrons, and neutrons to form different “atoms.” They should learn that if the number of protons is changed, the identity of the element is changed. If the number of neutrons is changed, the mass number is changed. If the number of electrons is changed, the charge is changed. (See Appendix A) Application/Analysis Level • Students use lima beans as a model of different isotopes. Students are given small beans and large beans. Students can calculate the average mass of each type of bean. Students are then given a sealed bean sample that contains some small and some large beans. Based on the average mass of their sample, they can calculate the percentage of each type of bean, using algebra. (See Appendix B) Synthesis/Evaluation Level • Vocabulary atom protons electrons neutrons charge atomic number average atomic mass mass number isotope periodic table half-life radioactive decay alpha radiation beta radiation gamma radiation Students can do research on radioactive isotopes to explore the following questions: o What factors influence the stability of a nuclide? o What are the different types of nuclear decay? o What happens when an atom undergoes nuclear decay? o Is there a way to predict what type of decay a particular radioisotope will undergo? o How can I use half-life data to calculate the amount of radioactive material present after a certain time period? SOL: CH.2 The student will investigate and understand that the placement of elements on the periodic table is a function of their atomic structure. The periodic table is a tool used for the investigations of (a) average atomic mass, mass number, and atomic number; (b) isotopes, half lives, and radioactive decay; (c) characteristics of subatomic particles as to mass and charge. August 1, 2009 SOL CH.2a,b,c – Appendix A PROTONS, NEUTRONS, AND ELECTRONS Name Part One – Basic Information about the Atom The ATOMIC NUMBER of Helium is 2. Atomic Mass Element Symbol This means that every atom of helium contains exactly 2 Atomic Number . Element Name The ATOMIC NUMBER can be used to identify an element. There are at least 110 different elements on the periodic table. Each element can be uniquely identified by its atomic number. Look at your periodic table. What is the ATOMIC NUMBER of Silver (Ag)? This means that each atom of silver contains PROTONS have a ELECTRONS have a . . charge, and they are located in the of the atom. charge, and they are located . The ATOMIC MASS of Helium is 4.00260. This number represents the average mass of all of the different isotopes of Helium that exist in nature. We will be discussing more information about isotopes later. For now, let us assume that we have an atom of helium that has a MASS NUMBER of 4. You should know that protons are much, much heavier than electrons. In fact, a proton is about 2000 times heavier than an electron. For this reason, scientists have assigned the following mass numbers to a proton and an electron. PROTON ELECTRON Mass Number 1 0 The ATOMIC NUMBER of a Helium atom is 2, which corresponds to 2 protons. If the MASS NUMBER of a Helium atom is 4, where are the other 2 mass units coming from? The answer is that this helium atom also contains 2 . Fill in the following table, based on the information for each of the three subatomic particles. Mass Number Charge PROTON ELECTRON NEUTRON Location in the Atom An atom is neutral and has no overall charge. This means that in the atom there should be the same number of and . The MASS NUMBER of an atom is equal to the sum of the and . Check with your teacher before proceeding to the next step. Part Two – Subatomic Particles Each of you should have a partner, and each of you should have a cup that contains the following items: PROTONS (brown beans) NEUTRONS (black beans) ELECTRONS (white beans) Look in your cup and count all of your beans. Write the numbers of each type of particle for your cup and for your partner’s cup in the table below. Your Cup PROTONS NEUTRONS ELECTRONS Your Partner’s Cup PROTONS NEUTRONS ELECTRONS Now we will review the symbol that chemists use to represent an atom. Let’s use helium-4 as an example: Mass Number (protons + neutrons) Atomic Number (protons) ! 2 He Based on the number of subatomic particles in your cup, write the symbol for the atom that is represented by these particles. Do the same thing for your partner’s cup. Your Atomic Symbol Your Partner’s Atomic Symbol Check with your teacher before proceeding to the next step. Part Three – Creating New Elements If you are the person with Fluorine, take ONE PROTON and ONE ELECTRON from your pile of beans and give them to your partner. Now fill in the table below with the new count of subatomic particles. Your Atom PROTONS NEUTRONS ELECTRONS Your Partner’s Atom PROTONS NEUTRONS ELECTRONS Did the chemical identity of your elements change when you made this change? If you said yes, then what are the NEW elements that you have now? If you said no, then explain why the chemical identity did NOT change. When an atom gains or loses a PROTON, what property of the atom is going to change? Based on the number of subatomic particles that you have now, write the symbol for the atom that you have. Write your partner’s symbol as well.. Your Atomic Symbol Your Partner’s Atomic Symbol Check with your teacher before proceeding to the next step. Part Four – Exploring Isotopes If you are the person with Oxygen, take TWO NEUTRONS from your pile and give them to your partner. Now fill in the table below with the new count of subatomic particles. Your Bag PROTONS NEUTRONS ELECTRONS Your Partner’s Bag PROTONS NEUTRONS ELECTRONS Did the chemical identity of your elements change when you exchanged NEUTRONS? If you said yes, then what are the NEW elements that you have now? If you said no, then explain why the chemical identity did NOT change. When an atom gains or loses a NEUTRON, what property of the atom is going to change? Based on the number of subatomic particles that you have now, write the symbol for the atom that you have. Write your partner’s symbol as well.. Your Atomic Symbol Your Partner’s Atomic Symbol Two atoms that have the SAME number of PROTONS and a DIFFERENT number of NEUTRONS are called . Are the atoms shown below are ISOTOPES of each other? Explain why or why not. Are the atoms shown below are ISOTOPES of each other? Explain why or why not. A chemical property is defined as “the ability of a substance to undergo a change that transforms it into a new substance.” For example, the ability of iron to rust is a chemical property. If you have 2 different isotopes of the same element, should they have the same chemical properties or different chemical properties? Explain why you think so. Check with your teacher before proceeding to the next step. Part Five – Charges and Ions If you are the person with Magnesium, take TWO ELECTRONS from your pile and give them to your partner. Now fill in the table below with the new count of subatomic particles. Your Bag PROTONS NEUTRONS ELECTRONS Your Partner’s Bag PROTONS NEUTRONS ELECTRONS Did the chemical identity of your elements change when you exchanged ELECTRONS? If you said yes, then what are the NEW elements that you have now? If you said no, then explain why the chemical identity did NOT change. When an atom gains or loses an ELECTRON, what property of the atom is going to change? A atom is neutral when the number of is equal to the number of An ION is defined as an atom that has a positive or negative charge. If you compare the number of protons and electrons, how can you determine the overall charge on an ION? If an atom LOSES ELECTRONS, then it will become charged. If an atom GAINS ELECTRONS, then it will become charged. The charge on an ion can be calculated as follows: (PROTONS) – (ELECTRONS) = charge on the ion Based on the number of subatomic particles that you have now, write the symbol for the atom that you have. Write your partner’s symbol as well.. Your Ion Symbol Your Partner’s Ion Symbol Check with your teacher to find out what to do next PLEASE MAKE SURE THAT THE CLEAR PLASTIC CUP HAS 9 BROWN BEANS, 9 WHITE BEANS, AND 10.BLACK BEANS. PLEASE MAKE SURE THAT THE RED PLASTIC CUP HAS 11 BROWN BEANS, 11 WHITE BEANS, AND 12.BLACK BEANS. PROTONS, NEUTRONS, AND ELECTRONS Name Summary Questions 1. When you change the number of PROTONS in an atom, you will change the 2. When you change the number of NEUTRONS in an atom, you will change the 3. When you change the number of ELECTRONS in an atom, you will change the 4. ISOTOPES are defined as two atoms that have the same number of but they each have a different number of 5. An ION is defined as an atom that has 6. True or False: “For a neutral atom, the number of protons should be equal to the number of electrons.” If you answered True, then explain WHY this statement is true. If you answered False, then write the complete symbol for a neutral atom in which the number of protons is NOT equal to the number of electrons. 7. Choose a neutral atom in which the number of protons is NOT equal to the number of neutrons. (Circle your answer) + -2 8. A student saw the following list of subatomic particles: 22P, 26N, 20E The student then wrote the atomic symbol as Is the student correct? -2 If yes, explain why. If not, then write the correct symbol 9. A student saw the following list of subatomic particles: 50P, 69N, 50E The student then wrote the atomic symbol as Is the student correct? If yes, explain why. If not, then write the correct symbol SOL CH.2a,b,c – Appendix B Exploring Isotopes of “Beanium” Name PRE-LAB QUESTIONS 63 65 1. Two different isotopes of copper ( Cu and Cu) are being compared to each other. For each property listed below, answer if you think that these two isotopes would have the SAME property or if you think that they would have DIFFERENT properties. (a) Number of PROTONS in the atom (c) Mass of the atom (b) Number of NEUTRONS in the atom (d) Chemical Reactivity 2. Suppose that your grades in Math class were based on Tests and Homework, and you received the following scores: Tests = 90 Homework = 80 (a) Your teacher tells you that he will count Tests as 50% of your grade and Homework as 50% of your grade. Calculate your Math Grade, based on these percentages (b) Suppose that your teacher changes his mind and decides that he will count Tests as 80% of your grade and Homework as 20% of your grade. Now calculate your new Math Grade, based on these new percentages. 3. At the end of the semester, a student’s grades were reported as follows: Nine Weeks Grade 90 Exam Grade 70 Semester Grade 85 (a) Based on the results shown above, do you think that the Nine Weeks Grade and the Exam Grade were each worth 50% of the Semester Grade? Answer Yes or No, and then explain why you think so. (b) Which grade (Nine Weeks or Exam) do you think counted for a higher percentage when determining the student’s Semester Grade? Explain why you think so. 10 11 4. There are two possible isotopes of Boron: B and B. Based on the average atomic mass found on the periodic table, which isotope of boron is more abundant in nature? Explain. Exploring Isotopes of “Beanium” Name Part One – Determining the Average Mass of a Single Atom of Beanium You should two cups of beans. One is labeled “Isotope A” and the other is “Isotope B.” 1. Place the small plastic cup labeled “A” on the electronic scale and press the zero button. The scale should read “0.00 g” with the cup on the pan. 2. Add the Beanium sample that is labeled “Isotope A” to the plastic cup. Record the mass of Isotope A in Data Table 1. 3. Count the individual “atoms” of Isotope A as you put them back in the larger cup. (There should be 20 “atoms.”) Record this number of atoms in Data Table 1. 4. Repeat Steps 1 through 3, using the small plastic cup labeled “B” and “Isotope B.” 5. Perform the calculations necessary to complete Data Table 1. (Hint: Average Mass of One Atom = Mass of Isotope divided by Number of Atoms) Data Table 1 Mass of Isotope A Number of Atoms in Isotope A Average Mass of One Atom of Isotope A Mass of Isotope B Number of Atoms in Isotope B Average Mass of One Atom of Isotope B Part Two – Determining the Average Mass for Different Samples of Beanium 1. Place TWO small plastic cups (one marked “A” and the other marked “B”) on the electronic scale and press the zero button. The scale should read “0.00 g” with the two cups on the pan. 2. Place 2 atoms of Isotope A into the A cup. Place 18 atoms of Isotope B into the B cup. 3. Record the Total Mass of the Beanium Sample in Data Table 2.. 4. Now adjust the number of atoms so that you have 5 atoms of A in the A cup, and 15 atoms of B in the B cup. Record the Total Mass of the Beanium Sample. 5. Now adjust the number of atoms so that you have 8 atoms of A in the A cup, and 12 atoms of B in the B cup. Record the Total Mass of the Beanium Sample. 6. Now adjust the number of atoms so that you have 12 atoms of A in the A cup, and 8 atoms of B in the B cup. Record the Total Mass of the Beanium Sample. 7. Now adjust the number of atoms so that you have 15 atoms of A in the A cup, and 5 atoms of B in the B cup. Record the Total Mass of the Beanium Sample. 8. Now adjust the number of atoms so that you have 18 atoms of A in the A cup, and 2 atoms of B in the B cup. Record the Total Mass of the Beanium Sample. Data Table 2 Number of Atoms of Isotope A Number of Atoms of Isotope B 2 18 5 15 8 12 12 8 15 5 18 2 Total Number of Atoms in the Beanium Sample Percentage of Isotope A Percentage of Isotope B Helpful Hints for Calculations Number of Atoms of Isotope A Percentage of Isotope A = -------------------------------------------- x 100% 20 Number of Atoms of Isotope B Percentage of Isotope B = ------------------------------------------ x 100% 20 Average Mass of One Atom of Beanium = Total Mass of Beanium Sample ÷ 20 Total Mass of Beanium Sample Average Mass of One Atom of Beanium Exploring Isotopes of “Beanium” Name ANALYSIS QUESTIONS 1. “Beanium” exists in two isotopes, which are called Isotope A and Isotope B. What is the main difference between Isotope A and Isotope B? 2. Chlorine exists in two isotopes, which are called chlorine-35 and chlorine-37. What is the main difference between chlorine-35 and chlorine-37? 3. Copy your data from Data Table 1 into the blanks below: Average Mass of One Atom of Isotope A: Average Mass of One Atom of Isotope B: 4. Suppose that you are given a sample of Beanium. You know that it contains 18 atoms of Isotope A and 2 atoms of Isotope B. You are asked to calculate the average mass of one atom of your sample. Should you use the following formula? Answer yes or no and explain why you think so. (Mass of Isotope A) + (Mass of Isotope B) -------------------------------------------------------- = Average Mass of Beanium 2 5. Calculate the average mass of Beanium for a sample that contains 90% Isotope A and 10% Isotope B. Use the numbers that you wrote down in Question 3. Show your work below, by filling in the blanks with the appropriate numbers. x Percent Abundance of A as a decimal = Average mass of Isotope A x Percent Abundance of B as a decimal = add these two values together Average mass of Isotope B Average mass of Beanium sample = 6. Compare your answer from Question 5 with the value from Data Table 2. (Look for the box that is heavily outlined in the lower right corner of Data Table 2.) Are these two values similar? Why? 35 37 7. There are two possible isotopes of Chlorine: Cl and Cl. Based on the average atomic mass found on the periodic table, which isotope of chlorine is more abundant in nature? Explain why you think so. 63 65 8. The element copper is present in two isotopes: Cu and Cu. The percent abundance of each isotope is listed below. Use these percentages to calculate the average atomic mass of copper. Isotope 63 Cu 65 Cu % Abundance 73% 27% Physical Science: Matter: Periodicity COURSE: Chemistry CONCEPT: SYSTEMS ENDURING UNDERSTANDING: A system is an organized collection of parts, which display observable and predictable patterns. Essential Understandings Assessment Samples – SOL/Blooms Vocabulary Students should understand: • Periodicity is regularly repeating patterns or trends in the chemical and physical properties of the elements arranged in the periodic table. • Horizontal rows called periods have predictable properties based on an increasing number of electrons in the outer orbitals. • Vertical columns called groups have similar properties because of their similar valence electron configurations. Knowledge/Comprehension Level • Students are given a periodic table and they should be able to identify the following features: periods, groups, alkali metal family, alkaline earth family, transition metals, halogens, and noble gases. Application/Analysis Level • Students should understand that if the formula of one compound is known, then the formulas of similar compounds are predictable. For example, if the formula of water is H2O, then the formulas of the other Group 16 hydrides should be H2S, H2Se, and H2Te. Synthesis/Evaluation Level • Students are given a set of cards that represent fictional elements. They must arrange them in order of increasing atomic mass (just as Dimitri Mendeleev did back in 1869). Students should look for similarities among the elements so they can be classified into groups/families. The cards will be set up so that there are two missing elements. These should create gaps in their table. Just like Mendeleev, students should be able to predict several properties of the missing elements, based on the existing trends. (See Appendix A) periodic law periods groups alkali metals alkaline earth metals transition metals halogens noble gases metals nonmetals metalloids (semimetals) atomic radius ionization energy electronegativity shielding effect electron configuration valence electrons SOL: CH.2 The student will investigate and understand that the placement of elements on the periodic table is a function of their atomic structure. The periodic table is a tool used for the investigations of (d) families or groups; (e) series and periods; (f) trends including atomic radii, electronegativity, shielding effect, and ionization energy. August 1, 2009 SOL CH.2def – Appendix A Physical Science: Matter: Periodicity COURSE: Chemistry CONCEPT: SYSTEMS ENDURING UNDERSTANDING: A system is an organized collection of parts, which display observable and predictable patterns. Essential Understandings Students should understand: • Electron configuration is the arrangement of electrons around the nucleus of an atom based on their energy level. • Atoms can gain, lose, or share electrons within the outer energy level. Assessment Samples – SOL/Blooms Vocabulary Knowledge/Comprehension Level • Students should be able to write the full electron configuration AND the noble gas abbreviated configuration for any element from hydrogen (Z=1) through krypton (Z=36). electron configuration orbitals (s, p, d, f) quantum numbers energy levels Aufbau Principle Pauli Exclusion Principle Hund’s Rule cation anion oxidation state Example: Titanium = 1s22s22p63s23p63d24s2 or [Ar] 3d24s2 Application/Analysis Level • Students should be able to recognize if a given electron energy diagram is correct or incorrect. If the diagram is incorrect, they should be able to identify which rule or principle (aufbau, Pauli, Hund) has been violated. (See Appendix A) Synthesis/Evaluation Level • Students should be able to give examples of several ions that are isoelectronic with a given noble gas. They should also be able to identify an element from its charge and configuration. Example: Name six ions that are isoelectronic with neon (N-3, O-2, F–, Na+, Mg+2, and Al+3) Example: A certain element, M, has the following formula for its oxide: M2O3. The M ion is isoelectronic with neon. What is the identify of element M? (Aluminum) SOL: CH.2 The student will investigate and understand that the placement of elements on the periodic table is a function of their atomic structure. The periodic table is a tool used for the investigations of (g) electron configurations, valence electrons, and oxidation numbers. August 1, 2009 SOL CH.2g – Appendix A Identify the rule (aufbau principle, Hund’s rule, or Pauli exclusion principle), if any, that has been violated in each of the following electron energy diagrams. If no rule has been violated, then state that it is correct as written. 3p 3s 3p 3s 2p 3p 3s 2p 3p 3s 2p 2p 2s 2s 2s 2s 1s 1s 1s 1s Physical Science: Matter: Structure COURSE: Chemistry CONCEPT: SCALE: Properties ENDURING UNDERSTANDING: Properties characterize objects, organisms, and substances. Essential Understandings Students should understand: • Matter is classified by its chemical and physical properties. • Physical properties refer to the condition or quality of a substance that can be observed or measured without changing the substance’s composition. • Chemical properties refer to the ability of a substance to undergo chemical reaction and form a new substance. Assessment Samples – SOL/Blooms Vocabulary Knowledge/Comprehension Level • Students should be able to classify a change as a physical change (e.g., ripping paper, melting ice) or a chemical change (e.g., burning paper, electrolysis of water) and explain their reasoning. Physical changes do not alter the chemical identity of the substances involved. Chemical changes involve a change in chemical identity and the formation of a new chemical substance. (See Appendix A) matter mixture pure substance homogeneous heterogeneous compound element physical property chemical property density conductivity melting point boiling point malleability ductility reactivity Application/Analysis Level • Students should be able to classify examples of matter as mixture vs. pure substance. If something is a mixture, students should be able to identify it as homogeneous or heterogeneous. If something is a pure substance, students should identify it as an element or a compound. If a mixture is separated or a compound is decomposed, students should know if that would be a physical or chemical change. (See Appendix B) Synthesis/Evaluation Level • Before discussing the concept of mixture and pure substance, students will discover the defining characteristics of these concepts on their own, through the concept-attainment model. Examples and non-examples are revealed to the students one at a time. o Examples (mixtures): coffee, tea, milk, gasoline, soda, steel, saltwater, blood, air o Non-examples (pure substances): gold, silver, iron, oxygen, nitrogen, water, ethanol, sucrose, salt, carbon dioxide Students will brainstorm to come up with ideas about what all the examples have in common. Students’ ideas will be evaluated continually as additional examples and non-examples are revealed. SOL: CH.2 The student will investigate and understand that the placement of elements on the periodic table is a function of their atomic structure. The periodic table is a tool used for the investigations of (h) chemical and physical properties. August 1, 2009 SOL CH.2h – Appendix A Laboratory Activity – Physical and Chemical Changes Part 1 – Signs of a Chemical Change There are FOUR basic observations that may suggest that a chemical reaction may have occurred. • • • • Change in color (out of nowhere) Formation of a gas Change in energy (heat, light, sound) Formation of a precipitate Add a few drops of phenolphthalein to the beaker that contains sodium hydroxide. 1. What do you observe? Now add a few drops of hydrochloric acid to the beaker until you see a color change. 2. What do you observe? Pour the contents of the beaker in the sink and rinse out the beaker with water. Add a piece of magnesium metal ribbon to the test tube that contains hydrochloric acid. Cover the test tube with another empty test tube (upside-down) to trap the gas. Hold the upside-down test tube in place. When all of the fizzing has stopped, call your teacher to put a flame to the gas. 3. What did you observe when the magnesium was added to the acid? 4. What did you observe when the gas in the upside-down test tube was lit? 5. Feel the outside of the test tube with the acid in it. How does it feel? Add about 3 drops of lead(II) nitrate to a clean, empty beaker. Now add about 3 drops of potassium iodide to the same beaker. 6. What do you observe? 7. All of the changes described in this section are classified as new substance was produced. changes because a Pour the contents of the beaker in the sink and rinse out the beaker with water. Part 2 – Phase changes 1. There are six phase changes that you need to memorize. They are listed below. Write the name of the phase change in each blank. solid ! liquid liquid ! solid liquid ! gas gas ! liquid solid ! gas gas ! solid 2. Look at the beaker with ice cubes and water in it. What do you observe on the outside of the beaker? 3. How do you suppose those drops of water got there? Now place the beaker on the hot plate. 4. What will happen to the ice cubes over time? 5. Once all of the ice has melted, what will happen to the water over time ? Add a piece of dry ice to a flask. Stretch a balloon over the flask. Place the flask on the hot plate. 6. What do you observe happening to the balloon? Your teacher will demonstrate phase changes involving iodine crystals. 7. When the iodine crystals change from a solid to a gas, this process is called 8. When the iodine vapor is changed from a gas to a solid, this process is called 9. All of the changes described in this section are phase changes. Phase changes are all classified as changes, because the chemical identity of the substance is . Part 3 – Physical or Chemical? For each of the following activities, you should check the boxes that apply, and then decide if the change is classified as PHYSICAL or CHEMICAL. Activity Take a piece of aluminum foil and rip it into smaller pieces. Add the aluminum pieces to a beaker, and then add some copper chloride solution to it. Squirt some methanol on the lab counter and wait for it to disappear. Squirt some methanol on the lab counter and call your teacher to put a flame near the liquid. Place an antacid tablet in the mortar and crush it with the pestle. Add some water to the powder in the mortar. Add about 3 drops of silver nitrate to a clean beaker. Then add about 3 drops of copper chloride to the beaker. TEACHER DEMO: Solid potassium chlorate is heated in a test tube until it turns into a liquid. TEACHER DEMO: A piece of candy is added to the molten potassium chlorate in the test tube. Did you see a color change? Did you see the formation of a gas? Did you notice a change in energy? Did you see the formation of a precipitate? Did you see a phase change? Is this a PHYSICAL or a CHEMICAL change? SOL CH.2h – Appendix B Laboratory Activity – Matter Part 1 – Classifying Matter At each lab station, you will see three samples of matter. Your job is to classify each sample as either a MIXTURE, a COMPOUND, or an ELEMENT. Each lab station has one of each type of matter represented. Station #1 Station #7 vinegar solution = copper sulfate salt (sodium chloride) = hydrogen peroxide solution copper = nitrogen Station #2 Station #8 sugar (sucrose) = oil and water aluminum = ammonium nitrate sedimentary rock = carbon Station #3 Station #9 zinc = hexane (C6H14) carbon dioxide = air bronze = helium Station #4 Station #10 potassium iodide = soda and spaghetti pieces pepper and water = oxygen hydrogen = cobalt chloride Station #5 Station #11 water = calcium carbonate calcium = green ink Gatorade = magnesium Station #6 Station #12 iodine crystals = copper baking soda (sodium bicarbonate) = methane (CH4) iron and sulfur = chicken noodle soup 1. A MIXTURE is defined as a sample of matter that represents a physical combination of more than one substance. Fill in the blanks below with the names of the samples of matter from each lab station that you classified as a MIXTURE: #1 #5 #9 #2 #6 #10 #3 #7 #11 #4 #8 # 12 Mixtures are classified as either HOMOGENEOUS or HETEROGENEOUS. What’s the difference between these terms? A HOMOGENEOUS MIXTURE appears to be uniform throughout. In other words, it appears to look like just one phase or one substance. On the other hand, in a HETEROGENEOUS MIXTURE you can see and identify more than one component or phase. 2. Sort all 12 mixtures into one of the following categories: HOMOGENEOUS MIXTURES 3. HETEROGENEOUS MIXTURES A COMPOUND is defined as a sample of matter that is formed from a chemical combination of more than one element. Fill in the blanks below with the names of the samples of matter from each lab station that you classified as a COMPOUND: #1 #5 #9 #2 #6 #10 #3 #7 #11 #4 #8 # 12 It is very important that you understand the difference between a COMPOUND and a MIXTURE. A COMPOUND has elements that are chemically bonded together. A MIXTURE has components that are only physically combined together. 4. If you wanted to separate a mixture into simpler substances, you could do it with a change. On the other hand, if you wanted to break down a compound into simpler substances, you would need to use a change. 5. An ELEMENT is defined as a sample of matter that CANNOT be broken down or decomposed by either a physical or a chemical change. Fill in the blanks below with the names of the samples of matter from each lab station that you classified as an ELEMENT: #1 #5 #9 #2 #6 #10 #3 #7 #11 #4 #8 # 12 6. You should realize that all of the substances listed above can be found on the Part 2 – Separation Techniques 1. Paper Chromatography You are going to be given a piece of paper with a line of green ink on it. You should lower the filter paper into the beaker of liquid so that only the bottom edge of the paper is touching the water. Over time, the water will rise up through the paper and separate the green ink into smears of blue and yellow. 2. a. Is the green ink from the marker classified as a compound or a mixture? b. In this experiment, are we producing a new substance that wasn’t there before? c. Is paper chromatography classified as a physical change or a chemical change? Filtration You should have a plastic cup with some water and pepper in it. You are going to use a piece of filter paper to separate the pepper from the water. a. Is pure water (H2O) classified as a compound or a mixture? b. Is the sample of water and pepper classified as a compound or a mixture? c. If you said “mixture” for question 2b., do think that water and pepper represents a homogeneous mixture or a heterogeneous mixture? How can you tell? Now you are going to pour the water and pepper mixture through a funnel lined with filter paper. d. Describe the appearance of the filter paper after all of the water has passed through it. e. Do you think the liquid that came through the filter paper should be classified as a compound or a mixture? Explain why. f. Filtration is a separation technique that is based on particle size. Do you think filtration should be classified as a physical change or a chemical change? 3 Distillation Your teacher will perform this experiment as a demonstration. a. Is solid copper sulfate (CuSO4) classified as a compound or a mixture? b. Is the combination of water and copper sulfate classified as a homogeneous mixture or a heterogeneous mixture? c. How can you tell? In distillation, the mixture is heated until it begins to boil. Which substance, water or copper sulfate, is going to appear in the collection flask first? d. Distillation is a technique used to separate two substances that have different e. When we use distillation to separate copper sulfate from water, are we producing a new substance that wasn’t there before? f. 4. Is distillation classified as a physical change or a chemical change? Electrolysis of water Your teacher will perform this experiment as a demonstration. a. Is pure water (H2O) classified as a compound or a mixture? b. During electrolysis of water, bubbles of gas are formed in two separate tubes. What two gases are being collected? and c. When we use electrolysis to break water down into its elements, are we producing new substances that weren’t there before? d. Is electrolysis classified as a physical change or a chemical change? e. The diagram at right shows a set-up for the electrolysis of water. The tube on the LEFT contains gas, and the tube on the RIGHT contains gas. Part 3 – Summary Questions 1. 2. Classify each of the following as COMPOUND, ELEMENT, HETEROGENEOUS MIXTURE, or HOMOGENEOUS MIXTURE green ink copper water carbon dioxide hydrogen bronze sand & water oxygen air saltwater sugar salt & sand Fill in the following “MATTER” chart with the words COMPOUND, ELEMENT, HETEROGENEOUS MIXTURE, HOMOGENEOUS MIXTURE, MATTER, and PURE SUBSTANCE. MIXTURE 3. Green ink can be separated easily by using 4. Sand & water can be separated easily by using 5. Saltwater can be separated easily by using 6. Water can be decomposed into hydrogen and oxygen by using 7. MIXTURES can be separated by changes. 8. COMPOUNDS can be separated by changes. 9. CANNOT be separated by either physical or chemical means. Physical Science: Matter: Structure COURSE: Chemistry CONCEPT: COMMUNICATION: Theory ENDURING UNDERSTANDING: Theories explain why natural phenomena occur, and they evolve to incorporate new knowledge. Essential Understandings Students should understand: • Discoveries and insights related to the atom’s structure have changed the model of the atom over time. • The modern atomic theory is called the Quantum Mechanical Model. Assessment Samples – SOL/Blooms Vocabulary Knowledge/Comprehension Level • Students can prepare a timeline that shows the important scientists who contributed to atomic theory, including a few details about their specific contribution. Scientists should include Democritus, John Dalton, J. J. Thomson, Ernest Rutherford, Robert Millikan, Niels Bohr, Dimitri Mendeleev, Henry Moseley, Max Planck, Werner Heisenberg, and Louis de Broglie. atom electron cathode ray tube nucleus planetary model periodic table quantum theory uncertainty principle particle-wave duality Application/Analysis Level • Students will observe a “think tube” which has various ropes extending from the sides. When a single rope is pulled as part of the teacher’s demonstration, students must draw a diagram of what they think the inside of the tube looks like. As more ropes are pulled, students are challenged to explain a few seemingly impossible results and then modify their diagrams accordingly. Students then can try to construct their own think tube that will behave similarly to the teacher’s model. Synthesis/Evaluation Level • Students can design their own mystery box, which contains a simple geometric shape (e.g., triangle, square, circle) that is sandwiched between two larger pieces of wood. The only way that the identity of the hidden shape can be revealed is that other students can shoot marbles at the mystery shape from different angles. Based on the deflection pattern of the marbles, students try to guess what shape is inside the box. This activity is somewhat parallel the Rutherford gold foil experiment. SOL: CH.2 The student will investigate and understand that the placement of elements on the periodic table is a function of their atomic structure. The periodic table is a tool used for the investigations of (i) historical and quantum models. August 1, 2009 Physical Science: Matter: Reactions COURSE: Chemistry CONCEPT: COMMUNICATION: Model ENDURING UNDERSTANDING: Models vary in complexity and facilitate understanding through the use of familiar concepts. Essential Understandings Students should understand: • Conservation of matter is represented in balanced chemical equations. • Chemical formulas are used to represent compounds. • Subscripts represent the relative number of each type of atom in a molecule or formula unit. • A coefficient is a quantity that precedes a reactant or product symbol or formula in a chemical equation and indicates the relative number of particles involved in the reaction. • Bonds form between atoms to achieve stability. Assessment Samples – SOL/Blooms Knowledge/Comprehension Level • Students should know the formula and charge of monoatomic ions (e.g., Na+, Mg+2, F–, O-2) based on their location on the periodic table. (See Appendix A) • Students should know the formula and charge of the following polyatomic ions: ammonium (NH4+), hydroxide (OH–), nitrate (NO3–), carbonate (CO32-), sulfate (SO42-), and phosphate (PO43-). Application/Analysis Level • Students should be able to write correct chemical formulas for molecular compounds, ionic compounds, and acids if they are given the name (and vice versa). They should also be able to identify the oxidation state of certain elements within a chemical formula. (See Appendix B) Synthesis/Evaluation Level • • • Students are given several chemical formulas that refer to a mystery element. They are also given information about the electron configuration of the mystery element when it becomes an ion. They must identify the element (See Appendix C) Students are given an incorrect Lewis dot structure. They must identify what is wrong with the structure and draw a correct Lewis structure. (See Appendix C) Students are given a molecular formula that contains a mystery element and the total number of valence electrons for the formula. They must draw the Lewis dot structure, identify the molecular geometry, and identify the mystery element. (See Appendix C) Vocabulary Law of Multiple Proportions empirical formula molecular formula structural formula covalent bond ionic bond ionization energy electronegativity polar nonpolar binary covalent binary ionic monoatomic ions polyatomic ions Lewis dot diagrams bonding pair nonbonding pair (lone pair) single bond double bond triple bond tetrahedral pyramidal bent trigonal planar linear SOL: CH.3 The student will investigate and understand how conservation of energy and matter is expressed in chemical formulas and balanced equations. (a) nomenclature; (b) balancing chemical equations; (c) writing chemical formulas (molecular, structural, empirical, and Lewis diagrams); (d) bonding types (ionic, covalent). August 1, 2009 SOL CH.3abcd – Appendix A At the top of each column, write the charge on the ion (oxidation state) that would be formed when these elements undergo chemical reactions. While the value for oxidation states can certainly vary, consider what the element would have to do to achieve a noble gas configuration. Your choices are +1, -1 +2, -2, +3, and -3 Common Oxidation States +4 H Li Na K Rb Cs Be Mg Ca Sr Ba B Al Ga In Tl C Si Ge Sn Pb 0 N P As Sb O S Se Te F Cl Br I He Ne Ar Kr Xe Rn SOL CH.3abcd – Appendix B 1. Binary Ionic Compounds sodium oxide NiBr3 magnesium chloride Ba3P2 iron(III) sulfide Al2O3 potassium nitride CoF2 calcium phosphide SrI2 2. Ionic Compounds with Polyatomic Ions (You need to memorize those polyatomic ions!) ammonium acetate Na2CO3 potassium chlorate Li2CrO4 magnesium hydrogen carbonate FeSO4 aluminum hydroxide Cu3(PO4)2 calcium nitrate Ba(CN)2 3. Binary Molecular Compounds carbon diselenide PBr5 sulfur trioxide SeF6 silicon tetrachloride ICl7 4. Acids hydrofluoric acid H2SO4 nitric acid HC2H3O2 5. Oxidation Numbers – Write the oxidation number of the element listed in bold type H2SO4 F2 N H3 SCl2 ClO3– PCl3 HBr PF5 HOCl Cr2O7-2 SOL CH.3abcd – Appendix C A certain element, X, forms the following ionic compounds: XCl2, XO, and X3N2 When an atom of element X becomes a stable ion, it has the electron configuration 1s22s22p63s23p6. a) Based on the given information, is element X classified as a metal or a nonmetal? b) When element X becomes a stable ion, what is the charge? (Don’t just say “positive” or “negative.” I want the sign and the number, like “+3” or “-2” c) What is the identity of element X? A student tried to write a Lewis dot structure for the molecule C3H3N. The student wrote a skeleton structure that was correct, but the arrangement of electrons was incorrect. The incorrect Lewis dot structure is shown on the left. a) How many valence electrons should be used for the molecule C3H3N? b) In the left box, circle all of the atoms (other than hydrogen) that do NOT have access to a full octet of electrons. c) In the right box, draw the correct Lewis dot structure for C3H3N. You should keep the same skeleton arrangement of atoms as the structure on the left. The number of electrons in your structure should be the same as the number you mentioned in part a). Make sure that all atoms (except hydrogen) have access to a full octet. INCORRECT Lewis Dot Structure CORRECT Lewis Dot Structure Consider the following information, as summarized in the table below. Note that the letters A, D, G, and J are used to represent elements from the periodic table. The letters O and F represent atoms of oxygen and fluorine. Molecule Total Number of Valence Electrons Additional Information AO2 16 none DO2 18 none GF3 24 JF3 26 The Lewis structure contains ALL single bonds The Lewis structure contains ALL single bonds (a) For each molecule listed in the table above, identify the molecular geometry/shape. You might want to sketch the Lewis electron dot structure on your own paper. AO2 = DO2 = GF3 = JF3 = (b) The molecules AO2, DO2, GF3, and JF3 represent covalent compounds. Based on the information in the table above, identify the elements A, D, G, and J. (More than one correct answer may be possible for each letter, but you only have to give ONE possible answer for each letter.) A= D= G= J= Physical Science: Matter: Reactions COURSE: Chemistry CONCEPT: CHANGE AND CONSTANCY: Cause and Effect ENDURING UNDERSTANDING: Observable changes occur in nature, and inferences can be made to explain their causes. Essential Understandings Students should understand: • Elements and compounds react in different ways. • Spontaneous reactions may be fast or slow. • Randomness (entropy), heat content (enthalpy), and temperature affect spontaneity. • Reaction rates/kinetics are affected by activation energy, catalysis, and the degree of randomness (entropy). Assessment Samples – SOL/Blooms Knowledge/Comprehension Level • Students should be able to balance an equation and identify the type of chemical reaction (See Appendix A) Application/Analysis Level • Students are given an equilibrium system. They should apply LeChatelier’s principle to predict the direction (if any) that the system will shift as a result of various changes to the system. For example: Consider an exothermic reaction. In which direction will the equilibrium position shift when the temperature of the system is increased? (Answer: to the left.) Synthesis/Evaluation Level • If students are given a sentence describing the chemicals that react together, they should be able to identify the type of reaction AND write a complete balanced chemical equation for that process. (See Appendix A) Vocabulary synthesis decomposition single replacement double replacement neutralization oxidation reduction endothermic exothermic equilibrium Le Chatelier’s Principle activation energy catalyst rate of reaction SOL: CH.3 The student will investigate and understand how conservation of energy and matter is expressed in chemical formulas and balanced equations. (e) reaction types (synthesis, decomposition, single and double replacement, oxidation-reduction, neutralization, exothermic and endothermic); (f) reaction rates and kinetics (activation energy, catalysis, degree of randomness). August 1, 2009 SOL CH.3ef – Appendix A Identify the TYPE of reaction that is represented in each diagram. Use the following choices: SYNTH DECOMP synthesis decomposition SR DR single replacement double replacement ! + 1) H + X M + ! + B OH C + M X C B + ! + 5) A H ! + 4) 6) OH combustion neutralization + ! 2) 3) COMB NEUTRAL D ! A + D BALANCE each reaction using the lowest possible whole numbers. Identify the TYPE of reaction. Use the following choices. SYNTH DECOMP synthesis decomposition SR DR single replacement double replacement 7) 8) 9) Fe(OH)3 Ca(C2H3O2)2 + ! (NH4)2CO3 ! COMB NEUTRAL combustion neutralization Fe2O3 + H2 O CaCO3 + NH4C2H3O2 Ni + Cl2 ! NiCl3 10) C3 H6 O + O2 ! CO2 + H2 O 11) F2 + ScBr3 ! ScF3 + Br2 + H2SO4 ! Al2(SO4)3 + H2 O 12) Al(OH)3 WRITE THE CORRECT FORMULAS for each substance listed. BALANCE each reaction using the lowest possible whole numbers. Identify the TYPE of reaction. Use the following choices. SYNTH DECOMP synthesis decomposition SR DR single replacement double replacement 13) octane (C8H18) + 14) diphosphorus pentoxide 15) hydrochloric acid 16) aluminum 17) potassium chlorate 18) iron(III) bromide + + oxygen + ! + carbon dioxide water magnesium hydroxide lead(II) nitrate ! ! COMB NEUTRAL ! ! water phosphoric acid (H3PO4) ! lead potassium chloride sodium sulfate + combustion neutralization magnesium chloride + + water aluminum nitrate + iron(III) sulfate oxygen + sodium bromide Physical Science: Matter: Structure COURSE: Chemistry CONCEPT: SCALE: Measurement ENDURING UNDERSTANDING: Measurement represents properties on a numerical scale. Essential Understandings Students should understand: • Atoms and molecules are too small to count by usual means. • A mole is a way of counting any type of particle (atoms, molecules, and formula units). • Stoichiometry involves quantitative relationships. • Stoichiometric relationships are based on mole quantities in a balanced equation. Assessment Samples – SOL/Blooms Vocabulary Knowledge/Comprehension Level • Students should know that there are 6.02 x 1023 particles in one mole. They should know that 1 mole of gas at standard temperature and pressure (STP) has a volume of 22.4 L. • Students should know how to use the periodic table to calculate the molar mass of a substance by using the atomic masses and adding up the total mass of all of the atoms that are found in the chemical formula. Avogadro’s number mole average atomic mass molar mass molar volume STP reactants products limiting reactant percent yield Application/Analysis Level • If given a balanced chemical reaction, students should be able to do the following conversions: mole-mole, mole-mass, mass-mass, and volumevolume. (See Appendix A) Synthesis/Evaluation Level • Students are given the amounts of each reactant in a chemical equation, and they are to do calculations to determine the limiting reactant, theoretical yield, and the percent yield (if given the actual yield). (See Appendix B) SOL: CH.4 The student will investigate and understand that quantities in a chemical reaction are based on molar relationships. (a) Avogadro's principle and molar volume; (b) stoichiometric relationships. August 1, 2009 SOL CH.4a,b – Appendix A Moles to Moles: KClO3 1 ! KCl + O2 (a) Balance the equation above. (b) If you started with 5.17 moles of potassium chlorate, how many moles of oxygen gas would be produced? Grams to Moles: Mg 2 + H3PO4 ! Mg3(PO4)2 + H2 (a) Balance the equation above. (b) How many moles of magnesium metal would be needed to react completely with 37.8 grams of phosphoric acid? Grams to Grams: C4H10 3 + O2 ! CO2 + H 2O (a) Balance the equation above. (b) How many grams of oxygen gas are required to react with 74.2 grams of butane (C4H10)? SOL CH.4a,b – Appendix B Limiting Reactant: K3PO4 + Ca(OH)2 ! KOH + Ca3(PO4)2 1. Balance the equation above. 2. Suppose that you reacted 41.5 grams of potassium phosphate with 32.4 grams of calcium hydroxide. (a) Determine which reactant is the limiting reactant, based on the product calcium phosphate. (b) Calculate the theoretical yield (in grams) of calcium phosphate. (c) When this experiment was performed (as described above), the mass of recovered calcium phosphate was equal to 24.7 grams. Calculate the percent yield for the reaction. Physical Science: Matter: Structure COURSE: Chemistry CONCEPT: SYSTEMS ENDURING UNDERSTANDING: A system is an organized collection of parts, which display observable and predictable patterns Essential Understandings Students should understand: • • • • • • • • Gases have mass and occupy space. Gas particles are in constant, rapid, random motion and exert pressure as they collide with the walls of their containers. Gas molecules with the lightest mass travel fastest. Relatively large distances separate gas particles from each other. An Ideal Gas does not exist, but this concept is used to model gas behavior. A Real Gas exists, has intermolecular forces and particle volume, and can change states. Equal volumes of gases at the same temperature and pressure contain an equal number of particles. Solutions can be a variety of solute/solvent combinations: gas/gas, gas/liquid, liquid/liquid, solid/liquid, gas/solid, liquid/solid, or solid/solid. Assessment Samples – SOL/Blooms Knowledge/Comprehension Level • Students should be able to write the basic postulates of the kinetic molecular theory. • Students should know the gas laws (Boyle, Charles, Gay-Lussac, ideal) and understand how to use them to solve for a particular variable. Application/Analysis Level • Students should be able to explain why each of the following events occur, in terms of the KMT. o A balloon will expand when heated. o Pressure inside a sealed can will increase when the can is heated. o A balloon will expand when taken to a higher elevation (at constant temp.) o Pressure inside a sealed can will increase when the can is squashed. Synthesis/Evaluation Level • • Students can perform an experiment that involves the calculation of the value of the ideal gas constant, R. (See Appendix A) Students can collect a sample of butane gas by water displacement. Based on the data collected, students can calculate the molar mass of the butane gas. Vocabulary pressure volume temperature kinetic-molecular theory Boyle’s Law Charles’ Law Gay-Lussac’s Law Dalton’s Law of partial pressures Ideal Gas Law ideal gases real gases solution concentration solute solvent molarity SOL: CH.4 The student will investigate and understand that quantities in a chemical reaction are based on molar relationships. (c) (d) (e) (f) partial pressure; gas laws; solution concentrations; chemical equilibrium. August 1, 2009 SOL CH.4c,d,e,f – Appendix A !"#!$#"%&'()%*+)&,+"#)("-)#".)!/'-%"'%) ! 012345614) ! 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I4!'/(!)..!#(!/+,)*-!!34!F!&7!3/4!52#!6!7!832!7/32+.!*(!)*+(!'3!-)+-2+)'(!'/(!?(%-(,'!(%%3%!43%!!! ! ! ! ! ! ! ! ! !!!! !!!!!!8)$!9! ! ! ! ! ! ! 832%!!"#!$%&!'()*+,)*-!!34!F#!</35!832%!-)+-2+)'&3,7#!J!)..!#(!/!K!!"#!$%&!'()*!J!L!"MMN! ! ! ! ! ! ! ! ! ! ! ! ! ! ! ! !!!)..!#(!/+ Physical Science: Matter: Structure COURSE: Chemistry CONCEPT: SCALE: Properties ENDURING UNDERSTANDING: Properties characterize objects, organisms, and substances. Essential Understandings Students should understand: • Two important classes of compounds are acids and bases. • Acids and bases are defined by several theories. • Acids and bases dissociate in varying degrees. Assessment Samples – SOL/Blooms Knowledge/Comprehension Level • Students can estimate the pH of various solutions (e.g., water, vinegar, ammonia, lemon juice, drain opener, etc.) by using colorful acid-base indicators and/or pH paper. Application/Analysis Level • Students can write balanced chemical equations for neutralization reactions. For example, hydrochloric acid in the stomach is neutralized with milk of magnesia (magnesium hydroxide). 2 HCl + Mg(OH)2 ! MgCl2 + 2 H2O Synthesis/Evaluation Level • Students can determine the molarity of household vinegar by performing a titration experiment with a standardized NaOH solution. (See Appendix A) Vocabulary Arrhenius definition Bronsted-Lowry definition acid base pH pOH molarity strong electrolyte weak electrolyte dissociation titration SOL: CH.4 The student will investigate and understand that quantities in a chemical reaction are based on molar relationships. (g) acid/base theory: strong electrolytes, weak electrolytes, and nonelectrolytes; dissociation and ionization; pH and pOH; and the titration process. August 1, 2009 SOL CH.4g – Appendix A Standardization of NaOH with KHP (potassium hydrogen phthalate, KHC8H4O4) Procedure: 1. Measure out approximately 1.00 gram of the solid powder KHP. Record the mass of KHP in the data table below. 2. Transfer the KHP into a 500 mL flask. Rinse the plastic weighing tray with distilled water to make sure that it is completely transferred. 3. Add approximately 100 mL of distilled water to the flask. 4. Add about 5 drops of phenolphthalein solution to the flask. 5. Make sure the buret is filled with NaOH solution so that the volume is in between 0 and 5 mL. 6. Record the exact volume of the buret to two decimal places in the data table below. 7. Place the flask directly under the buret. Slowly add the NaOH and swirl to mix it. 8. As the pink color starts to persist, try to add the NaOH drop by drop. 9. When one drop of NaOH results in a pink color that does not fade, you have reached the equivalence point in the titration. 10. Record the final volume on the buret to two decimal places. 11. Rinse out your flask several times with with water. You don’t have to dry it, but it should be clean. 12. Repeat Steps 1-11. You may need to add more NaOH to the buret in step 5. Trial 1 Trial 2 mass of KHP (3 sig figs) g molar mass of KHP (KHC8H4O4) g/mol moles of KHP (3 sig figs) mol mol initial volume of NaOH in buret (two decimal places) mL mL final volume of NaOH in buret (two decimal places) mL mL volume of NaOH used (two decimal places) mL mL moles of NaOH used mol mol liters of NaOH used L L molarity of NaOH mol/L mol/L Average value for Molarity of NaOH solution g mol/L DETERMINING THE MOLARITY OF ACETIC ACID IN VINEGAR In this experiment you will be titrating a vinegar solution with the sodium hydroxide that you have standardized. Vinegar contains acetic acid (CH3CO2H). Sodium hydroxide is a base. You will be using phenolphthalein indicator to help you visualize the endpoint. When all of the acid has been neutralized by the base, the color of the solution will turn pink. You will do 2 trials. Based on your results, you can calculate the molarity of the acetic acid. Write a balanced equation for the neutralization of acetic acid with sodium hydroxide: Trial 1 volume of vinegar (3 sig figs) 10.0 Trial 2 mL 10.0 mL initial volume of NaOH in buret (two decimal places) mL mL final volume of NaOH in buret (two decimal places) mL mL volume of NaOH used (two decimal places) mL mL liters of NaOH used L L Molarity of NaOH (from earlier experiment) M M moles of NaOH mol mol moles of vinegar L L molarity of vinegar mol/L mol/L Average value for Molarity of vinegar solution mol/L What is the percentage of acetic acid by mass in your vinegar sample? Let’s assume that the density of vinegar is close to 1 g/mL. Plug in your molarity into the calculations below and find out. mol CH3CO2H x L 60 g CH3CO2H x 1L x 1 mol CH3CO2H 1000 mL 1 mL 1g x 100% = Physical Science: Matter: Structure COURSE: Chemistry CONCEPT: COMMUNICATION: Theory ENDURING UNDERSTANDING: Theories explain why natural phenomena occur, and they evolve to incorporate new knowledge. Essential Understandings Students should understand: • Atoms and molecules are in constant motion. • The Kinetic Molecular Theory is a model for predicting and explaining gas behavior. • Forces of attraction between molecules determine the physical changes of state. • Vapor pressure is a property of a substance determined by intermolecular forces. Assessment Samples – SOL/Blooms Knowledge/Comprehension Level • If students are given a vapor pressure diagram for a certain liquid, they should be able to determine the normal boiling point of the liquid. Application/Analysis Level • • If students are given a vapor pressure diagram for several liquids, they should be able to predict which liquid has the strongest/weakest intermolecular forces. (See Appendix A) If students are given a vapor pressure diagram, they should be able to identify the boiling point for conditions when the pressure is reduced below atmospheric pressure. (See Appendix A) Vocabulary kinetic-molecular theory pressure temperature volume vapor pressure intermolecular forces London dispersion forces dipole-dipole forces hydrogen bonding Synthesis/Evaluation Level • Students should be able to analyze the structural formulas of various substances and identify the type of intermolecular forces that each substance has. They should predict which substance in a given pair has a higher boiling point. (See Appendix B) SOL: CH.5 The student will investigate and understand that the phases of matter are explained by kinetic theory and forces of attraction between particles. (a) pressure, temperature, and volume; (b) vapor pressure. August 1, 2009 SOL CH.5ab – Appendix A Why is zinc a good conductor of electricity, but solid zinc chloride is not a good conductor of electricity. Explain your answer by discussing the properties of each solid at the particle level. Match these substances with their boiling points using your knowledge of intermolecular forces. o magnesium fluoride A. –188 C fluorine B. hydrogen fluoride C. 1400 C o 20 C o The diagram below shows how the temperature of a substance changes as it is being heated. Label each portion of the graph with either a phase of matter or a phase change Identify the letter that matches each of the following: 1.00 atm B D C Pressur e E solid liquid vapor A triple point F G critical point normal MP Temperatur e normal BP Based on the phase diagram below, answer the following questions: 2.00 atm 1.00 atm ` o a. At –10 C and 0.50 atm, what phase of matter would this substance be in? b. If the temperature were to then increase to o 30 C at constant pressure, what phase change would occur? c. If the pressure were to then increase to 1.5 atm at constant temperature, what phase change would occur? 0.50 atm d. If the temperature were to then o decrease to 0 C at constant pressure, what phase change would occur? e. From the graph, estimate the normal BP of this substance: -20 -10 0 10 20 30 40 50 60 70 80 90 Temp (oC) Based on the vapor pressure curves above, answer the following questions: a. Which liquid has the highest vapor pressure at room temperature? b. Which liquid has the highest boiling point? c. Which liquid experiences the weakest intermolecular forces? d. If only one of these liquids experiences hydrogen bonding, which one would it be? o e. If liquid B is at a temperature of 40 C, what would the equilibrium vapor pressure be? f. What is the normal boiling point of liquid C? SOL CH.5ab – Appendix B For each substance below, indicate whether the substance is polar or nonpolar. Then indicate the type of intermolecular forces that is most appropriate. Use the following abbreviations: LDF = London dispersion forces Substance Polar or Nonpolar? H2 HCl H 2O HF Type of IMF DD = dipole-dipole forces Substance Polar or Nonpolar? CO CO2 CH3OH CH3F Type of IMF HB = hydrogen bonding Substance Ne NH3 PH3 C 3H 8 For each pair of substances below, circle the one that should have the higher boiling point: a.) Ne b.) CH4 c.) N2 Xe C4H10 CO d.) HF e.) NH3 HBr PH3 Polar or Nonpolar? Type of IMF Physical Science: Force, Motion, and Energy: Energy COURSE: Chemistry CONCEPT: SCALE: Properties ENDURING UNDERSTANDING: Properties characterize objects, organisms, and substances. Essential Understandings Students should understand: • Solid, liquid, and gas phases of a substance have different energy content. • Specific amounts of energy are absorbed or released during phase changes. • Specific heat capacity is a property of a substance. • Polar substances dissolve ionic or polar substances; nonpolar substances dissolve nonpolar substances. • The number of solute particles changes the freezing point and boiling point of a pure substance. Assessment Samples – SOL/Blooms Vocabulary Knowledge/Comprehension Level • Students should know that likes dissolve likes. For example, they should know that water is polar and that polar solutes will dissolve in water, whereas nonpolar solutes do not dissolve in water. Likewise, two nonpolar substances (e.g., turpentine and gasoline) should be very miscible with each other. solid liquid gas specific heat capacity polar nonpolar freezing point boiling point heat of fusion heat of vaporization phase diagram vapor pressure triple point heating curve solute solvent Application/Analysis Level • • • Students can do calorimetric calculations to calculate the amount of heat required to raise the temperature of a certain mass of water by a certain number of degrees Celsius. (See Appendix A) If given the molar heat of fusion (or molar heat of vaporization) for a certain substance, they can calculate the amount of heat required to melt (or evaporate) a certain quantity of that substance. (See Appendix A) Students can calculate the molar heat of fusion for ice (See Appendix B) Synthesis/Evaluation Level • Students can record the mass of water in a Styrofoam cup. Then they can dissolve a known mass of an ionic solid into the water and record the change in temperature as the solute dissolves completely. Based on the temperature change, the mass of the water, and the mass of the solute, they can calculate the molar enthalpy of solution. SOL: CH.5 The student will investigate and understand that the phases of matter are explained by kinetic theory and forces of attraction between particles. (c) phase changes; (d) molar heats of fusion and vaporization; (e) specific heat capacity; (f) colligative properties. August 1, 2009 SOL CH.5cdef – Appendix A 38. The molar heat of fusion for ice is 6.02 kJ/mol. Use this information to answer the following questions. a. How much heat energy is required to melt 125 grams of ice? b. How many grams of ice at 0 °C can be melted by the addition of 1500 J of heat? SOL CH.5cdef – Appendix B CALCULATING THE MOLAR HEAT OF FUSION OF ICE Introduction Water freezes at 0oC. Ice melts at the same temperature, 0oC. The amount of heat energy required to melt one mole of ice at its melting point is known as the molar heat of fusion. The units for the heat of fusion are normally given as kilojoules per mole (kJ/mol). Most students are more familiar with the calorie (cal) than they are with the joule (J). One calorie is defined as the amount of heat required to raise the temperature of 1 gram of water by 1 degree Celsius. The conversion factor between calories and joules is as follows. 1 cal = 4.18 J In this experiment you will melt a small quantity of ice in a Styrofoam cup that contains warm water. You will then be able to use your data to calculate the molar heat of fusion of ice. Laboratory Procedure 1. There should be a beaker filled with water at your lab station, as well as two Styrofoam cups nested together. Record the mass of the clean, dry cups to the nearest 0.01 g. 2. Carefully pour water into the Styrofoam cup until the cup is approximately one-third full. Record the mass of the cups with the water to the nearest 0.01 g. 3. Place the Styrofoam cup in a 400 mL beaker. This is done so that the cup does not tip over. Record the initial temperature of the warm water to the nearest 0.1oC. 4. Obtain a cup of ice cubes and place them on a paper towel. Blot the excess water from them so they are dry. IMMEDIATELY after you have recorded the water temperature, carefully add two or three ice cubes to the water. The total mass is important, so do not allow the water to splash or spill out of the cup. 5. Gently stir the ice water with the thermometer. Do not touch the bottom of your cup with the thermometer. Your goal is to reach a temperature of exactly 0oC. 6. Once the water temperature reaches 0oC, place the thermometer on the counter so that it doesn’t roll off the table. Carefully remove any extra pieces of ice with the plastic spoon. Try to remove ice only (not water). 7. Record the mass of the Styrofoam cups with water and melted ice to the nearest 0.01 g. 8. Pour the water into the sink and place the cups upside-down on a towel to dry. Return any unused ice to your teacher. Dry off the countertop. 9. If time permits, you can repeat the entire experiment to get data for a second trial. Recorded Data Trial 1 Trial 2 Mass of two clean, dry Styrofoam cups (to nearest 0.01 g) g g Mass of cups + water (to nearest 0.01 g) g g Initial temperature of water (to nearest 0.1 oC) o Mass of cups + water + melted ice (to nearest 0.01 g) C g o C g Calculations (Use the data for either Trial 1 or Trial 2) 1. Calculate the initial mass of water used in this experiment. Show your calculations below. 2. Calculate the amount of heat (in joules) which was transferred from the water to the ice in this experiment. Use the following equation, and show your work below. heat = (mass of water, in grams) x 4.18 J g oC x (change in temperature, in oC) 3. Calculate the mass of ice that melted in this experiment. Show your calculations below. 4. Calculate the moles of ice that melted in this experiment. Show your calculations below. 5. Calculate the heat of fusion of ice, in units of kilojoules per mole. (Round answer to 2 sig figs.) Show your calculations below. 6. The accepted value for the molar heat of fusion of ice is 6.0 kJ/mol. Calculate your percent error, based on your answer in #5.