2021 AP Chemistry Summer Assignment: Review & New Material for First Year & AP Chemistry, Slides of Chemistry

Download 2021 AP Chemistry Summer Assignment: Review & New Material for First Year & AP Chemistry and more Slides Chemistry in PDF only on Docsity! AP Chemistry 2021 Summer Assignment Dear AP Chemistry Students, The purpose of this summer packet is to provide you with a complete review of first year Chemistry (with some remediation work necessitated by the pandemic) and to begin new material that will be covered in the course. Before starting this assignment, please first print it out (please try to double side or take to a print shop so they can do it for you! Let’s save some hydrocarbons!). To complete this assignment, you will watch a number of different videos (QR codes accompany each section) while completing notes. You will then complete practice problems on your own to ensure that you understand the concepts. This packet is lengthy as last year’s AP Chemistry students firmly believed that a complete review of first year Chemistry and introduction to AP material over the summer makes the transition much easier. Do not wait to complete this assignment until the end of the summer (as you will note there are days during the summer for us to meet on Teams to go over this work) and please make sure to email me if you have any questions or concerns (Rebecca.Poliner@providencehigh.org). You will have a test after the first day of class. These are the STRONGLY suggested virtual review meetings (will last as long as you need). After the third meeting, I will give you a Summer Assignment Practice Test which will look VERY similar to your test. The night before (9 PM) each meeting, you are expected to scan and email your completed work to Ms. Poliner, even if you are unable to attend. Please use THIS LINK (or email me for it) to join our Teams meetings. -June 23 @ 2 PMà Pg. 1-22 . -July 14 @ 2 PM à Pg. 23-44 (not including Le Châtelier’s Principle). -August 4 @ 2PMà Pg. 44-60 If you fail the Summer Assignment Test, we will have a meeting to discuss whether AP Chemistry is the best match for you. This is not meant as punishment, rather to ensure that every AP Chemistry student succeeds in this course. Best, Ms. Poliner 1 Comments from Previous Students: What does a student need to do to succeed in this class? • I believe the students should devote a lot of time and effort into AP Chemistry. They should not go into the class with the mindset that the class will be easy. They have to constantly practice, and they should go outside of class to get extra help if need be. They should be willing to complete all the work assigned on time because if you fall behind, you will struggle. Furthermore, they should work efficiently, especially if they have other AP classes. • I would recommend this class if you find yourself to be hardworking and would not be afraid to ask questions. This is one of the keys to success in AP Chemistry and many other rigorous courses. I learned how to ask for help in this class and came in to practice a lot. • I would highly recommend the class especially if going into a STEM field. It is difficult at times but if you ask questions and come for help you will succeed. Make sure to fully understand the questions and topics before moving on and do not rush. • To succeed, you have to be willing to sacrifice a lot of your time to ask for help and complete assignments. I would recommend the class because you truly learn a lot and you’ll feel good about yourself when you finish the course. • You need to commit a majority of your school time to this class. If you don’t get a concept, do practice problems until you do. • Make sure to watch the videos and don’t copy from other people because, in the end, you have to review all of the information of your quizzes. Take the class seriously and Ms. Poliner is your best friend. She helps you with everything so don’t rely on classmates too much. This class forces you to talk to your teacher. • You’re going to put in just as much time outside of the class as inside the class, probably more. 4 Practice: 1. I want to complete a lab in which I need to weigh 0.540 g of a substance. I use a balance that can measure to the nearest 0.01 g. Will I be able to determine the mass of this substance to three significant figures? If not, how many will I be able to measure it to? 2. I want to complete a lab in which I need to weigh 0.540 g of a substance. I use a balance that can measure to the nearest 0.001 g. Will I be able to determine the mass of this substance to three significant figures? If not, how many will I be able to measure it to? Mole Conversions: Write the conversion factors on each line: Mass (g) Mole Particles (atoms, molecules, formula units) Volume (L) (remember that the substance must be a ____________________) Ex. 0.5 mol of C6H12O6 to grams Ex. How many molecules of hydrogen are in 3.36 L of hydrogen gas at STP? Ex. 45.6 g of Na2SO4 to formula units. 5 Ex. What mass of iron can be recovered from 25.0 g of Fe2O3? Practice: 1. Convert 35 g of CuSO4 • 5 H2O to moles. Hint: The 5 molecules of water should be considered part of this molecule (Ans: 0.14 mol) 2. 4.3 x 1024 molecules of carbon dioxide to liters at STP (Ans: 160 L) 3. What mass of silver can be produced from 125 g of Ag2S? (Ans: 109 g) Net Ionic Reactions: States of Matter: Solubility Rules that you must memorize: 6 Ex: 3Cu(NO3)2 (aq) + 2Na3PO4 (aq) à Cu3(PO4)2 (s) + 6NaNO3 (aq) Ex2. Cl2 (g) + 2KBr (aq) à 2KCl (aq) + Br2 (g) Practice: Write the net ionic reaction for each of the following. 1. 2AgNO3 (aq) + Na2CrO4 (aq) à Ag2CrO4 (s) + 2NaNO3 (aq) 2. Zn (s) + CuSO4 (aq) à ZnSO4 (aq) + Cu (s) 9 Concentration Calculations: Molarity (M) Concentration: Ex. Calculate the molarity of a solution that has 0.127 mol of sucrose in 655 mL of solution. Practice: 1. Calculate the molarity of 1.60 L of a solution containing 1.55 g of dissolved KBr. (Ans: 0.00814 M) Creating a Solution from a Solid: Ex. What mass of glucose is needed to make 105 mL of a 1.02M glucose (C6H12O6) solution? 1. I need 100. mL of a 0.20 M solution of Mg(OH)2. I only have a solid. What mass of magnesium hydroxide do I need to weigh to make this solution? (Ans: 1.2 g) 10 Creating a Solution by Dilution: Ex. What volume of 5.5 M stock KCl solution is needed to make 2.5 L of a 0.100M KCl solution? 1. I need 150 mL of a 0.50 M solution of HCl. I have a 6.0 M stock solution. What volume of stock solution would I need to make this new, dilute solution? (Ans: 12.5 mL) 11 Practice: For the two example problems and both of the practice problems above (pg. 9-10), clearly describe the steps necessary to make each of the solutions. Make sure to reference which of the tools above you would use for each measurement. First Ex: First #1: Second Ex: Second #1: 14 Gas Pressure: Ex. 1277 mmHg to atm Practice: Convert the following. 1. 11.5 kPa to mmHg (Ans: 86.3 mmHg) 2. 0.558 atm to mmHg (424 mmHg) Relationships between measurements of gases: o Temperature and Pressure: o Temperature and Volume: o Volume and Pressure: Units of Pressure: 15 The Ideal-Gas Law: • Ideal Gas Law: Ex. If I have 4 moles of a gas at a pressure of 4,256 mmHg and a volume of 12 liters, what is the temperature? Practice: 1. If I have 3 moles of gas in a container with a volume of 60 liters and at a temperature of 400K, what is the pressure inside the container? (Ans: 2 atm) 2. Determine the number of grams of carbon dioxide in a 450.6 mL tank at 1.80 atm and -50.5 °C. (Ans: 1.95 g) 3. Fritz Haber, a German chemist, discovered a way to synthesize ammonia gas (NH3) by combining hydrogen and nitrogen gases at extremely high temperatures and pressures. a) Write the balanced equation for this reaction. b) If 18.2 kg of nitrogen combines with excess hydrogen at 580°C and 240 atm, what volume of ammonia gas is produced? (Ans: 380 L) 16 Molar Mass of a Gas: Ex1. What is the molar mass, in grams per mole, of a gas if a 3.16 g sample of gas at 0.750 atm and 45°C occupies a volume of 2.05L? Ex2. What is the molar mass of a gas with a density of 3.50 g/L at a temperature of 345 K and a pressure of 0.950 atm? Practice: 1. If 9.006 grams of a gas are enclosed in a 50.00 liter vessel at 273.15 K and 2.000 atmospheres of pressure, what is the molar mass of the gas? What gas is this if it is diatomic? (Ans: 2.02 g/mol; H2) 2. What is the density (in g/L) of a gas with molar mass 130 g/mol at 0.8 atm and 29 °C? (Ans: 4 g/L) 3. The density of an unknown gas at 10°C and 739 mm Hg is 1.31 g/L. Calculate the molar mass of the gas. (Ans: 31.3 g/mol) 19 2. A 410. mL sample of O2 is collected over H2O at 25°C and 759.0 mm Hg pressure. How many moles of oxygen gas is present? (Vapor pressure of water at 25°C = 23.8 mm Hg) (Ans: 0.0162 mol) 3. A reaction occurs that produces a gas. After being collected over water, the gas has a mass of 1.211 g and occupies a volume of 677 mL. The reaction occurs at 23°C and an atmospheric pressure of 1.02 atm. Calculate the molar mass of the gas. (Vapor pressure of water at 25°C = 23.8 mm Hg) (Ans: 44.0 g/mol) Oxidation Numbers: Assigning Oxidation Numbers: 1. A way of keeping track of electrons gained and electrons lost. 2. Complete the following rules for assigning oxidation numbers: a. Elements always have an oxidation number of _________. b. The sum of the oxidation numbers of compounds is ________. c. The sum of the oxidation numbers for polyatomic ions is ______________________________________________. d. Oxygen is normally _________ except when it is a peroxide (O2 2-) in which case the oxidation number is ___________. e. Hydrogen is ________ when bound to nonmetals and ______________ when bonded to metals. f. Fluorine is always _______. Other halogens typically have an oxidation number of ______ but when combined with oxygen in molecules, they have ________________________ oxidation states. g. Group 1 Metals are always _______. Group 2 Metals are always_______. h. Positive Charge generally comes first. 20 Ex. Find the oxidation states for the following compound: a. KMnO4 b. SO42- Practice: Assign oxidation numbers to the following substances. KNO3 Al(NO3)3 NH3 O2 MgH2 H2O2 H2SO3 CO32- Na2O2 OF2 Redox Reactions: 1. How do we know if a reaction is a redox reaction? 2. If an element in a reaction is oxidized, why must there also be an element that is reduced? Section 4.4- Oxidation-Reduction Reactions Identifying Oxidation and Reduction 1. Reduction means _____________________________________. Oxidation means _____________________________________. 2. If an element is oxidized, what happens to the charge? __________ What happens if an element gains electrons (reduced)? __________ Practice: In the reactions below, note what is oxidized and what is reduced. Explain how you know. Ex1. CI4 + 2 Br2 à CBr4 + 2 I2 Ex2. 2H2O2à2H2O + O2 21 Practice: 1. Li + FeCl3 à 3 LiCl + Fe 2. HNO3 + NaHCO3 à H2O + CO2 + NaNO3 Subatomic Particles: The Modern View of Atomic Structure: • Protons (positively charged) and neutrons (neutrally charged) in the center, electrons (negatively charged) around the outside. • Because the size of an atom is so small, rather than using grams, we use amu (atomic mass unit). • Isotopes- same number of protons, different number of neutrons (have same chemical properties but different physical properties) • Mass Number= # protons + # neutrons • Atomic Number = # protons • C à Mass Number-12; Atomic Number- 6. Practice: Ion Number of Protons Number of Electrons Number of Neutrons Mass Number N3- 14 19 18 39 Br- 45 3 2 4 24 Empirical Formulas from Analyses: • Empirical Formula: Exa. A compound with elements C, H, and O is found to have 9.1% hydrogen and 54.5% carbon. What is the empirical formula? Exb. A compound with an empirical formula of C2H4O has a molecular mass of 176 g/mol. What is the molecular formula? *Tips* 1. Carry decimal places out at least 4 digits. 2. You can only round to a whole # if it is within 1 tenth of that number (Ex. 1.03à 1) 3. When you are not within 1 tenth you will need to multiply all of your answers so they are whole numbers. (Ex. 1.5 x 2 = 3; 1.2 x 5 = 6) Practice: 1. A compound is found to be 64.9 % carbon, 13.5% hydrogen, and 21.6% oxygen. Its molecular mass is 148 g/mol. What is its molecular formula? (Ans: C8H20O2) 25 Determining Empirical Formula Using Combustion Reactions: This is content that is normally taught in Honors Chemistry but was not last year due to us being BOLD. Please make sure you pay extra attention hereJ Ex. Answer the following questions about a pure compound that contains only carbon, hydrogen, and oxygen. A 0.7579 g sample of the compound burns in O2(g) to produce 1.9061 g of CO2(g) and 0.3370 g of H2O. a. Calculate the individual masses of C, H, and O in the 0.7579 g sample. b. Determine the empirical formula Practice: 1. 12.915 g of a biochemical substance containing only carbon, hydrogen, and oxygen was burned in an atmosphere of excess oxygen. Subsequent analysis of the gaseous result yielded 18.942 g carbon dioxide and 7.749 g of water. Determine the empirical formula of the substance. (Ans: CH2O) 26 2. Vitamin C is an organic acid containing only carbon, hydrogen and oxygen. A 1.00 g sample of vitamin C is analyzed by combustion. The compound produced 1.50 g of carbon dioxide and 0.41 g of water. If the molar mass of Vitamin C is 176.12 g/mol, what is the molecular formula of Vitamin C? (Ans: C6H8O6) Wavelength, Frequency, and Energy: The Wave Nature of Light: • All types of electromagnetic radiation move through a vacuum at 3.00 x 108m/s (speed of light). Ex. A certain microwave has a wavelength of 0.032 meters. Calculate the frequency of this microwave. Practice: 1. A radio station broadcasts at a frequency of 590 KHz. What is the wavelength of the radio waves in nanometers? (Ans: 5.1 x 1011nm) 29 Orbitals: 1. How many orbitals are in the 2s sublevel? 2. How many orbitals are in the 2p sublevel? 3. How many orbitals are in the 3p sublevel? 4. What is the difference between the 2s and 2p sublevel? 5. What is the difference between the 2p and 3p sublevel? Ex. O*: Practice: For the following atoms, write the shorthand electron configuration for elements that are underlined and the longhand electron configuration AND the orbital diagram for the elements that are starred. 1. N: 2. Co*: 30 3. Fe2+*: 4. O2-: Lewis Dot Diagrams: Ex1. Zn: 31 Drawing Lewis Dot Diagrams: Steps: Cà Sà Là Cà Ex. H2O Correct Lewis Dot Diagram (do not draw the version without the VSEPR model) Ex2. PBr3 Ex3. O2 Ex4. SO2 VSEPR Model: Model used in chemistry to predict the shape that molecules assume in order to minimize electron repulsions. 34 Periodic Trends: Moving Across the Periodic Table: • Atomic Radius decreases because… • Ionization Energy increases because… • Electronegativity increases because… Moving Down the Periodic Table: • Atomic Radius increases because… • Ionization energy decreases because… • Electronegativity decreases because… 35 Practice: 1. Which has a larger atomic radius, Na or Mg? Explain. 2. Which has a larger ionization energy, Br or I? Explain. 3. Which has a smaller electronegativity, Al or S? Explain. 4. Which has a smaller atomic radius, F or Cl? Explain. 5. Which has a smaller ionization energy, O or F? Explain. 6. Which has a larger electronegativity, Cl or I? Explain. 7. Which has a larger atomic radius, S or N? Explain. 8. Which has a greater ionization energy, Li or Na? Explain. 36 Polarity of Molecules and Attractive Forces: Electronegativity: Bond Polarity: 39 Ex3. CH3OH versus CH3CH2CH2CH2CH2CH2OH Ex4. HOCH2CH2CH2OH versus HOCH2CH2CH3 Practice: 1. This graph shows the BP’s of analogous compounds using elements from periods 2, 3, 4, and 5. a. Explain why the BP of Xe > Kr > Ar > Ne: b. Why is the BP of H2O > the others in its group? 2. Why is ΔHvap (energy required to vaporize a substance) much greater than ΔHfus (energy to melt a substance)? What does this reveal concerning changes in intermolecular forces in going from solid to liquid to vapor? 3. For which molecule in each of the following pairs would you expect the stronger intermolecular forces? Make sure to identify the type of intermolecular force EACH experience. a. CH3CH2CH2NH2 or H2NCH2CH2NH2 b. CH3CH3 or H2CO 40 c. CH3OH or H2CO d. HF or HBr Equilibrium Constants: This is content that is normally taught in Honors Chemistry but was not last year due to us being BOLD. Please make sure you pay extra attention hereJ Chemical Equilibrium: What are the three factors that affect reaction rate? Make sure to explain each of them. 1. 2. 3. 41 Equilibrium Constants: As always, we can express the position of equilibrium numerically: aA + bB ßà cC + dD Ex. Determine the value of the equilibrium constant for the reaction below if [N2O4]=1.5x10-3M and [NO2]=0.571M. N2O4(g) ⇌ 2 NO2 (g) Practice: 1. Write the equilibrium expression for each of the following: a. 2CO(g) + O2(g) ⇌ 2CO2(g) b. NH4NO3(s) ⇌ N2O(g) + 2H2O(l) 2. For the reaction N2 (g) + 3H2 (g) ⇌2NH3 (g): a. Write the equilibrium expression below. b. Find the Keq for the reaction if [N2]= 0.230 M, [H2] = 0.120 M, and [NH3] = 0.133 M. c. Using the Keq from “b”, determine the [NH3] if [N2] = 3.0 x 10-2 M and [H2] = 3.7 x 10-2 M. 44 3. Given the following equilibrium constant at 427ºC, Na2O (s) ⇌ 2 Na (l) + ½ O2 (g) K1 = 2 x 10-25 NaO (g) ⇌ Na (l) + ½ O2 (g) K2 = 2 x 10-5 Na2O2 (s) ⇌ 2 Na (l) + O2 (g) K3 = 5 x 10-29 NaO2 (s) ⇌ Na (l) + O2 (g) K4 = 3 x 10-14 Determine the values for the equilibrium constant for the following reactions: a. 2Na2O2 (s) ⇌ O2 (g) + 2Na2O (g) (Ans: 6 x 10-8) b. NaO (g) + Na2O (s) ⇌ Na2O2 (s) + Na (l) (Ans: 8. 10-2) Le Chatelier’s Principle: This is content that is normally taught in Honors Chemistry but was not last year due to us being BOLD. Please make sure you pay extra attention hereJ Le Châtelier’s Principle: 1. Effect of Concentration Changes on Equilibrium: H2CO3 ßà CO2 + H2O Add H2CO3- Add CO2- Remove H2CO3- • How can we change concentration? 45 2. Effect of Pressure/Volume Change on Equilibrium: N2 (g) + 3H2 (g) à 2NH3 (g) Increase Pressure- Decrease Pressure- 3. Effect of a Change in Temperature on Equilibrium 2SO2 + O2 ßà2SO3 + heat Add Heat- Remove Heat- • How to determine where “heat” is in the reaction: • Does the equilibrium constant (Keq) change when a shift occurs? 46 Practice: 1. Consider the following equilibrium for which ΔH<0 2 SO2(g) + O2(g) ßà 2 SO3(g) 
How will each of the following changes affect an equilibrium mixture of the three gases (shift right, shift left, no change) AND how will Keq change? (a) O2(g) is added to the system (b) the reaction mixture is heated; (c) the volume of the reaction vessel is doubled (d) a catalyst is added to the mixture (e) the total pressure of the system is increased by adding a noble gas (f) SO3(g) is removed from the system 2. Given the exothermic reaction below: 2NO (g) + H2 (g) ßàN2O (g) + H2O (g) In what direction would the reaction shift to overcome each of the following stresses: a. Adding H2:___________________________ b. Removing N2O:___________________________ c. Increasing pressure:___________________________ d. Decreasing volume:___________________________ e. Adding a catalyst:___________________________ f. Increasing temperature:___________________________ 3. Given the endothermic reaction below: NaCl (s) ßà Na+ (aq) + Cl- (aq) If I add AgNO3, in what direction would the reaction shift, and why? 49 5. What is the conjugate acid of each of the following bases? a. CN- ___________ b. SO4-2 ___________ c. H2O ___________ d. HCO3- ___________ The pH Scale: pH – “Power of Hydronium” measures the pH = - log [H+] pOH=-log[OH-] pH + pOH = 14 Calculate the pH and the pOH of the following: Ex. [H+] = 1 x 10-3 M pH = pOH = 1. [H+] = 1 x 10-9 M pH = pOH = 2. [H+] = 1.2 x 10-5 M pH = pOH = 3. [H+] = 2.3 x 10-1 M pH = pOH = Ex. [OH-] = pH= 2.0 pOH= Practice: [H+] pH [OH-] pOH Acidic/Basic 1. 5.6 x 10-5 M 2.0 2. 2.50 3. 3.4 x 10-3 M 4. 2.0 x 10-3 M 5. 6.25 6. 5.6 x 10-2 M 50 New AP Chemistry Material: Molecular Geometry: (This video is a mix of first year material and new material) Note: These molecular geometries must be memorized for the first day of school! Section 9.2- The VSEPR Model • Electron Domain Geometry: o Bonding Domain: o Nonbonding Domain: ar) 5a] barry om: Be eee lg Electron Domains around a Central Atom Number of _Electron- Electron Domain Bonding Nonbonding Molecular Domains Geometry Domains Domains Geometry Example 5 5 0 PCls Trigonal Trigonal bipyramidal bipyramidal 4 1 SF, Seesaw 3 2 a CIF; T-shaped Linear 6 6 0 SEs Octahedral Octahedral 5 1 BrF5 Square pyramidal 4 2 XeFs Square planar 51  54 Mass Spectrometry: 1. What does each peak represent? 2. Calculate the atomic weight of copper. 1. The diagram below represents the spectrum of chlorine, consisting of five peaks, labeled I, II, III, IV, and V, respectively. Peak I is due to the *“CI° ion. Relative abundance 30 35 40 45 SO 55 60 65 70 75 80 Mass / charge a. What analytical technique would give a spectrum like that shown above? b. How many different isotopes exist for chlorine atoms? c. State why the spectrum of chlorine consists of more peaks than your answer above. d. Suggest what peaks II and IV might be due to. 55  56 ______1. Which statement is true regarding the relative abundances of the 6lithium or 7lithium isotopes? A) The relative proportions change as neutrons move between the nuclei B) 7Lithium is much more abundant C) The relative ratio depends on the temperature of the element D) 6Lithium is much more abundant ______2. All isotopes of an element possess the same: A) number of electrons, atomic number and mass, but have nothing else in common B) atomic number and mass, but have nothing else in common C) chemical properties and mass, but have nothing else in common D) number of electrons, atomic number and chemical properties ______3. 63Cu is 69% of the naturally occurring isotope of Cu. If only one other isotope is present for natural copper, what is it? A) 59Cu B) 65Cu C) 61Cu D) 62Cu 4. Identify each element based on the mass spectra below: a. _________________________________ b._________________________________ c. Explain the spectra below: 59 Practice: 1. Identify the element based on the PES above. 2. Identify the element based on the PES below. 60 Ex. If an element has a greater ionization energy, all of its peaks will be the ________________________________ of the element that has the lower ionization energy. Practice: 1. Draw the PES for Na below. Then, on the same diagram, draw the PES for Mg in a different color. Make sure to clearly label each. 2. Draw the PES for Al below. Then, on the same diagram, draw the PES for Mg in a different color. Make sure to clearly label each.