AP Chemistry Summer Homework 2018 Welcome to AP Chem!!!! In early May next year, you will sit for the AP exam in chemistry. This exam covers the content of an introductory college-level chemistry course. To be successful on this exam and in the course, you will be expected to do at least 1 hour of homework each night during the school year (including weekends and vacations). When we meet in the fall, you will need to have a working knowledge of basic chemistry concepts. To help prepare you I have developed the following guidelines for summer work: 1. You are expected to complete three assignments over the summer. Each assignment has a reading component, a note-taking component, and a problem-set. 2. As you read each chapter, outline important information. Outlines should include definitions of all bold-faced terms and answers to all sample exercises within the reading (show all your work!!!). Outlines will be part of your AP Chemistry notebook. I will leave it up to you to organize your own notebook any way that works for you however an organized notebook is essential. Within each chapter you will find interesting, real-world applications of chemical concepts developed in the text. These Chemical Impact segments are required reading and will be discussed in class. 3. The problem sets for each assignment (found in the boxes below) must be mailed to my home address by their due date: Mrs. Furlong 24 Bears Den Way Columbia, CT 06237 4. Each problem set is worth 100 homework points. As part of each set, I have assigned black problems and blue problems. Blue problems have the answers in the back of the book. It is your responsibility to check your blue answers for accuracy. If your answer is not correct go back and work on the problem for a reasonable period of time. If you still cannot figure it out circle the number and plan to stay for extra help the first or second day of school. 5. SHOW YOUR WORK!!!! If you do not show your work, you will not receive any credit for the problem. Remember to label all numbers with units. 6. As you complete your summer homework, keep a log of the time it takes you to read/outline/and answer the problem sets. I will be collecting this information on the first day of school. Please feel free to contact me with any questions you have (e-mail: nfurl413@colchesterct.org). I will do my best to help in any way I can. We have a long and challenging road ahead of us so work hard (and have a great summer!!!) See you in September.
Assignment #1 (6/22 7/1) Read/outline chapter 1 Chemical Foundations Hint: Complete Review questions as you read/outline. Outlines should include the following: scientific method theory vs. natural law SI system of units (Tables 1.1, 1.2) mass vs. weight Why is there uncertainty in all measurements? precision vs. accuracy significant digits/calculations with sig. digits - see attached worksheet o Memorize Rules for Counting Significant Digits o Memorize Rules for Significant Digits in Calculations o Memorize Rules for Rounding dimensional analysis: upstairs/downstairs disappears - see attached worksheet temperature scales (Celsius, Kelvin. Fahrenheit) o Memorize how to convert between the Celsius and Kelvin scales o (K = C + 273) density Compare/contrast the 3 main states of matter Fig. 1.16 Classification of matter o Pure substance vs. mixture o Heterogeneous mixture vs. homogeneous mixture o Element vs. compound separation techniques: filtration, distillation, chromatography o Is filtration an appropriate method to separate a homogeneous mixture (solution)? physical change vs. chemical change
ChemQuest 2 Significant Figures Scientific notation can be a very nice way of getting rid of unnecessary zeros in a number. For example, consider how convenient it is to write the following numbers: 32,450,000,000,000,000,000,000,000,000,000,000 = 3.245 x 10 34 0.000000000000000127 = 1.27 x 10-16 There are a whole lot of zeros in the above numbers that are not really needed. As another example, consider the affect of changing units: 21,500 meters = 21.5 kilometers 0.00582 meters = 5.82 millimeters Notice that the zeros in 21,500 meters and in 0.00582 meters are not really needed when the units change. Taking these examples into account, we can introduce three general rules: 1. Zeros at the beginning of a number are never significant (important). 2. Zeros at the end of a number are not significant unless (you ll find out later) 3. Zeros that are between two nonzero numbers are always significant. Therefore, the number 21,500 has three significant figures: only three of the digits are important the two, the one, and the five. The number 10,210 has four significant figures because only the zero at the end is considered not significant. All of the digits in the number 10,005 are significant because the zeros are in between two nonzero numbers (Rule #3). Critical Thinking Questions 1.Verify that each of the following numbers contains four significant figures. Circle the digits that are significant. a) 0.00004182 b) 494,100,000 c) 32,010,000,000 d) 0.00003002 2. How many significant figures are in each of the following numbers? a) 0.000015045 b) 4,600,000 c) 2406 d) 0.000005 e) 0.0300001 f) 12,000
I nformation : The Exception to Rule #2 There is one exception to the second rule. Consider the following measured values: It is 1200 miles from my town to Atlanta. It is 1200.0 miles from my town to Atlanta. The quantity 1200.0 miles is more precise than 1200 miles. The decimal point in the quantity 1200.0 miles means they used a more precise tool and measured the distance right down to a tenth of a mile. So.Rule #2: Zeros at the end of a number are not significant unless there is a decimal point in the number. A decimal point anywhere in the number makes zeros at the end of a number significant. Critical Thinking Questions 3. Verify that each of the following numbers contains five significant figures. Circle the digits that are significant. a) 0.00030200 b) 200.00 c) 2300.0 d) 0.000032000 4. How many significant figures are there in each of the following numbers? a) 0.000201000 b) 23,001,000 c) 0.0300 d) 24,000,410 e) 2400.100 f) 0.000021 In numerical problems, it is often necessary to round numbers to the appropriate number of significant figures. Consider the following examples in which each number is rounded so that each of them contains 4 significant figures. Study each example and make sure you understand why they were rounded as they were: 42,008,000 is rounded to 42,010,000 12,562,425,217 is rounded to 12,560,000,000 0.00017837901 is rounded to 0.0001784 120 is rounded to 120.0
5. Round the following numbers so that they contain 3 significant figures. a) 173,792 b) 0.0025021 c) 0.0003192 d) 30 6. Round the following numbers so that they contain 4 significant figures. a) 249,441 b) 0.00250122 c) 12,049,002 d) 0.00200210 Information : Multiplying and Dividing When you divide 456 by 13 you get 35.0769230769 How should we round such a number? The concept of significant figures has the answer. When multiplying and dividing numbers, you need to round your answers to the correct number of significant figures. To round correctly, follow these simple steps: 1) Count the number of significant figures in each number. 2) Round your answer to the least number of significant figures. So. 456 has 3 sig dig and 13 has 2 sig dig. Our answer can only have 2 sig dig. We have to round our answer above to 2 sig dig: 35.0769230769 is rounded to 35 Critical Thinking Questions 7. Solve the following problems. Make sure your answers are in the correct number of significant figures. a) (12.470)(270) = b) 36,000/1245 = c) (310.0)(12) = d) 129.6/3 = e) (125)(1.4452) = f) 6000/2.53 = 8. Round the following numbers to the tens place. a) 134,123,018 = b) 23,190.109 = c) 439.1931 = d) 2948.2 =
Information : Adding and Subtracting Did you know that 30,000 plus 1 does not always equal 30,001? In fact, usually 30,000 + 1 = 30,000! I know you are finding this hard to believe, but let me explain Recall that zeros in a number are not always important, or significant. Knowing this makes a big difference in how we add and subtract. For example, consider a swimming pool that can hold 30,000 gallons of water. If I fill the pool to the maximum fill line and then go and fill an empty one gallon milk jug with water and add it to the pool, do I then have exactly 30,001 gallons of water in the pool? Of course not. I had approximately 30,000 gallons before and after I added the additional gallon because 30,000 gallons is not a very precise measurement. So we see that sometimes 30,000 + 1 = 30,000! Rounding numbers when adding and subtracting is different from multiplying and dividing. In adding and subtracting you round to the least specific decimal place of any number in the problem. ( It s all about the decimal place!!!! ) Example #1: Adding 10.6 + 1.007 = 11.607 Since 10.6 goes to 1 decimal place and 11.607 goes to 3 decimal places...our answer can only go to 1 decimal which is 11.6. Example #2: Subtracting 95830-18.99 = 95811.01 Since 95820 has 4 sig dig (the zero is NOT significant) this number is significant to the tens place. 18.99 is significant to the hundredths place. So..our answer can only go to the tens place which is 95810 Critical Thinking Questions 9. a) 24.28 + 12.5 = b) 120,000 + 420 = c) 140,100 1422 = d) 2.24 0.4101 = e) 12,470 + 2200.44 = f) 450 12.8 =
10. The following are problems involving multiplication, dividing, adding, and subtracting. Be careful of the different rules you need to follow! a) 245.4/120 = b) 12,310 + 23.5 = c) (31,900)(4) = d) (320.0)(145,712) = e) 1420 34 = f) 4129 + 200 = Dimensional Analysis Worksheet Upstairs/downstairs disappears You are required to KNOW the metric system as well as 1 mole = 6.02 x 10 23 molecules 1 cm 3 = 1 ml U se dimensional analysis to solve the following problems. SHOW ALL YOUR WORK!! NO WORK NO CREDIT!! 1. One-step conversions: Ex: 10.2 mm x 1 cm = 1.02 cm 10 mm a. 107 cm to m b. 30.0 g to kg c. 6.5 ml to L d. 2.5 moles to molecules e. 5.92 x 10 24 molecules to moles 2. Two-step conversions: Ex: 0.0576 L x 1000 ml x 1 cm 3 = 57.6 cm 3 1 L 1 ml a. 63 cm to km b. 1.56 x 10 7 mg to kg
c. 2800 cm 3 to L d. Extra Credit Challenge!!!! 0.348 L to mm 3 Assignment #2 (7/1 7/31) Read/outline chapter 2 Atoms, Molecules, and Ions Outlines should include the following: Identify the contributions of scientists in the early history of chemistry. Law of conservation of mass, Law of definite proportion, Law of multiple proportions Dalton s atomic theory Avogadro s hypothesis JJ Thomson and cathode ray tube experiments o What evidence led him to believe the electron is negative? o Why did he assume electrons are present in all elements? Millikan and oil drop experiment Rutherford and gold foil experiment How has picture of atom changed over time? Compare/contrast subatomic particles based on relative charge, mass and location within the atom atomic number, mass number, isotopes molecule vs. ion, cation vs. anion covalent bond vs. ionic bond Formation of ions (Fig. 2.22 VERY IMPORTANT): cations = +, loss of 1 or more electrons anions = -, gain of 1 or more electrons polyatomic ions (Make flash cards for table 2.5: name on one side and formula, including charge, on other side. Memorize polyatomic ions.) Periodic table Groups(or families): alkali metals, alkaline earth metals, chalcogens, halogens, noble gases Metals vs. nonmetals vs. metalloids Memorize naming rules: A. ionic compounds cation 1 st, anion 2 nd cations a. Roman numerals indicate charge (oxidation state) of transition metals b. o us and i c rule (hint: o us = l o wer charge, i c = h i gher charge) c. ium = polyatomic Anions a. Monatomic end in ide b. Polyatomic ate and ite ; define oxyanion
B. acids ide to hydro ic ate to ic ite to ous C. molecular (nonmetals only) use table 2.6 Problem set #2 Due date mailed by: 7/31 Review Questions (pg. 68, 69) # s 2, 5, 6, 8, 9 Active Learning Questions (pg. 69) #s 11, 12, 13 Problems: 23, 27, 28, 30, 35-39, 46, 48, 51, 53, 54, 72, 73, 74, 81 Assignment #3 (8/2 8/22) Read/outline chapter 3 Stoichiometry (sects. 3-1 to 3-8) Outlines should include the following: average atomic mass THE MOLE!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!!! 1 mole = 6.02 x 10 23 particles = grams (fill in molar mass) calculating molar mass - see attached worksheet Amu vs. molar mass % composition Empirical vs. molecular formula Calculating empirical formula Chemical equations understand what a coefficient represents What is conserved in an equation? Coefficient vs. subscript Balance chemical equations (Law of conservation of mass) - see attached worksheet
Calculating Molar Mass Worksheet Molar mass is the total mass in grams of one mole of a compound. To find it, you add the average atomic masses (found on the periodic table of elements) of the all the atoms in the compound Example: Mg(NO 3 ) 2 Mg = 24.31 x 1 = 24.31 grams N = 14.01 x 2 = 28.02 grams O = 16.00 x 6 = 96.00 grams 148.33 grams 1. Explain where the numbers 1, 2, and 6 came from in the example above. 2. Find the molar mass of O 2 3. Find the molar mass of Sr(OH) 2 4. Find the molar mass of Al 2 (SO 4 ) 3
5. Find the molar mass of C 6 H 12 O 6 6. Find the molar mass of Ca 3 (PO 4 ) 2 Worksheet: Balancing Equations Fill in the blanks with the most appropriate term: A tells the story of a chemical reaction. are the starting substances in the reaction while are the new substances that are formed. The large numbers in front of some of the formulas are called. These numbers are used to the equation because chemical reactions must obey the Law of of Matter. The number of atoms of each element on both sides of the equation must be because matter cannot be or. When balancing equations, the only numbers that can be changed are ; remember that must never be changed in order to balance an equation. II. Balance the following equations: 1. Al + O 2 Al 2 O 3 2. C 3 H 8 + O 2 CO 2 + H 2 O 3. Al(NO 3 ) 3 + NaOH Al(OH) 3 + NaNO 3 4. KNO 3 KNO 2 + O 2 5. O 2 + CS 2 CO 2 + SO 2
6. KClO 3 KCl + O 2 7. BaF 2 + K 3 PO 4 Ba 3 (PO 4 ) 2 + KF 8. H 2 SO 4 + Mg(NO 3 ) 2 MgSO 4 + HNO 3 9. Al + H 2 SO 4 Al 2 (SO 4 ) 3 + H 2 10. WO 3 + H2 W + H 2 O