AP Chemistry Summer Packet 2020-2021 PDF

Dear AP Chemistry Student, I’m excited that you are thinking about taking AP Chemistry in the 2020-2021 school year with me. This is a hard class that requires dedication, a fair amount of work and a love for chemistry. This year’s course may be structured a little differently, or we might have some...

Dear AP Chemistry Student, I’m excited that you are thinking about taking AP Chemistry in the 2020-2021 school year with me. This is a hard class that requires dedication, a fair amount of work and a love for chemistry. This year’s course may be structured a little differently, or we might have some new challenges given the global situation, but we will get through it! AP Chemistry is meant to be a second-year course….that is, there is a lot of new material to cover and very little time to go over topics we studied in Honors Chemistry or Chemistry I. However, because Chemistry is comprehensive, we can’t forget about all the concepts learned in the last school year. To that end, I’ve put together this notes and problems packet for you to complete during the summer. Don’t try to do it all at once – but practicing a little at a time every few days will help to flex those brain muscles over the summer. This packet will form the first chapter of your AP Chemistry Notebook. At the end of the first week of class, I will evaluate your completion of this packet and the first assessment in AP Chemistry will follow shortly thereafter (within the first three weeks of class). If you make an effort to review this material, I don’t anticipate you having any difficulty with the first test. In that spirit, enjoy your summer and I look forward to a new year of AP Chemistry with you. Cheers, Dr. K  GASES, LIQUIDS, AND SOLUTIONS P = pressure V = volume PV = nRT T = temperature moles A n = number of moles P A = Ptotal x X A , where X A = ----,----------:-- total moles 111= mass P I olal = P A + P B + Pc +... M= molar mass D= density In n = M KE= kinetic energy v= velocity K = °C + 273 A= absorbance D= m a = molar absorpti vity V b = path length KE per molecule = i mv 2 c = concentration Molarity , M = moles of solute per liter of solution Gas constant, R = 8.314 J mol-I K- 1 = 0.08206 L atm mol-I K- 1 A = abc = 62.36 L torr mol-I K- 1 ) atm = 760 mm Hg = 760 torr STP = 273. 15 K and 1.0 atm Ideal gas at STP = 22.4 L mol-I THERMODYNAMICS I ELECTROCHEMISTRY q = heat q = mc!'lT m = mas s c = specific heat capacity !'lSO = L.S D products - L.S D reactants T = temperature S° = standard entropy MiO = ~Mi7 products - ~MiJ reactants HO = standard enthalpy GO = standard Gibbs free energy !'lGO = ~!'lGJ products - ~!'lGJ reactants n = number of moles EO= standard reduction potentia) I = current (amperes) = -RTln K ' q = charge (coulombs) t = time (seconds) I = !1t Faraday's constant, F = 96,485 coulombs per mole of electrons I joule 1 volt = I coulomb -4- ,  AP Chemistry Summer Review Part I: Physical & Chemical Changes, Matter & Energy 1. Label each as either physical or chemical change. a. corrosion of aluminum metal by hydrochloric acid b. melting wax c. pulverizing an aspirin tablet d. digesting a Three Musketeers bar e. explosion of nitroglycerin f. a burning match g. metal warming up, due to the burning match h. water vapor condensing on the metal i. the metal oxidizes, becoming dull and brittle j. salt being dissolved by water 2. For each process described, state whether the material being discussed (in bold) is a mixture or compound, and state whether the change is physical or chemical. a. An orange liquid is distilled (boiled to separate components with different boiling points), resulting in the collection of a red solid and a yellow liquid. b. A colorless, crystalline solid is decomposed, leaving a pale yellow-green gas and and a soft, shiny metal. c. A cup of tea becomes sweeter as sugar is added to it. 3. Classify each as mixture (homogeneous or heterogeneous) or pure substance (elements or compounds). a. water b. blood c. the oceans  d. iron e. brass (an alloy of zinc and copper) f. wine g. sodium bicarbonate (baking soda) 4. Explain how the five states of matter and energy are related. (HINT: Think of the motion of the particles!) 5. Consider the burning of gasoline and the evaporation of gasoline. Which represents a physical change and represents a chemical change? Give the reason for your answer. 6. A) Label the arrows on the diagram below with the correct phase change processes. B) Draw a particle diagram representing each phase. Solid Liquid Gas 7. Describe the three main intermolecular forces and explain how their relationship is important in determining a compound’s state of matter at a particular temperature. → This is a major concept on the AP Chem Exam! AP Chemistry Summer Review Part II: Uncertainty in Measurement and Calculations: 1. Exact Numbers: Counted numbers and definitions do not involve any measurement and are considered as exact numbers Definitions: 1 week = 7 days. 1 mile = 5,280 feet 1 yard = 3 feet Counted: 5 Players on the basketball court. 23 students in a room 25 pennies used by a class in an experiment. 5 rocks 2. Measured Numbers: All measured numbers have some degree of uncertainty. When recording measurements, record only the significant figures. Record measurements to include one decimal estimate beyond the smallest increment on the measuring device. Examples (consider a measuring instrument like a ruler): If smallest increment = 1m, then record measurement o 0.1m (i.e. 3.1m) If smallest increment = 0.1m, then record measurement to 0.01m (i.e. 5.67 m) If smallest increment = 0.01m, then record measurement to 0.001m (i.e.12.675 m) c. Unless otherwise stated the uncertainty in the last significant figure (the uncertain digit) is assumed to be ±1 unit. Modern digital instruments and many types of volumetric glassware will state the level of uncertainty. 3. Rules for counting Significant Figures. a.Non-Zero Numbers: Non-zero numbers are always significant. b. Zeros: 1: Leading zeros that come before the first non-zero number are never significant 2. Captive zeros (sandwich zeros) that fall between two non-zero digits are always significant. 3. Ending zeros that appear after the last non-zero digit are significant only when a decimal point appears somewhere in the number. Examples: Number 0.005 5005 5005.00 500. 0.0050 Sig Figs 1 4 6 3 2 c. Scientific Notation: Significant figures are recorded in the mantissa (number 1 ≤ x < 10) Examples: 3 5 -23 4 Number 3.0 x 10 5.998 x 10 6.00000 x 10 0.5 x 10 Sig Figs 2 4 6 1 4. Rules for Using Significant Figures in Calculations (a) Multiplication, Division, Powers and Roots:-“LEAST SIG.FIG RULE” 1. The result should be reported to the same number of significant figures as the measured number having the least number of significant figures. 2. Only consider the number of significant figures in each of the measured numbers! (not constants) (b) Addition and Subtraction: “LEAST PRECISE DECIMAL RULE” 1. The result should be reported with the same decimal precision as the measured number having the uncertain digit in the least precise decimal place. 2. Only consider the decimal precision in each of the measured numbers! (not constants) (c) Addition/Subtraction combined with Multiplication/Division 1. Always perform the addition portion of the calculation 1st to determine the correct decimal precision of the sum. (least precise decimal rule) 2. Once the precision of the sum has been determined you can count the number of significant figures in the sum to apply the “least sig.fig rule” in performing the multiplication. 3. Do not round until the final calculation has been completed. (d) Scientific Notation with different powers of 10:  Problems How many significant figures in the following numbers: 1. ______ 1,245m 2. _______ 0.030m 23 3. _______ 10,000m 4._______ 1.340 x 10 m 14 5. _______ 3.02003 x 10 m 6. _______ 0.0000001m 7. _______ 1,000. 8. _______ 0.10000010 9: Convert the following numbers into standard scientific notation: 4 a. 96.3 x 10 g ____________________ 23 b. 0.05 x 10 s ____________________ -7 c. 123 x 10 m ____________________ Problems 10 – 18: Perform the following Calculations and record your answers in the proper number of significant figures and units. 10. 0.6030s + 0.82s = 11. 4.1m + 0.3789m – 153.22m = 2 4 6 12. 3.1567 x 10 g + 9.212 x 10 g – 4.677 x 10 g = 0.307𝑔 13. (1.0× 10−3 )𝑚𝑙 = 1.26×10−3 𝑘𝑔 14. (3.2𝑚 + 10𝑚 + 8.9𝑚)(4.3×10−6 𝑠) = 3 15. √5.33 × 105 𝑚 =  Part II: Simple Metric Conversions and Consistent Units Section 1: Metric Conversions One of the major benefits of using the metric system is the ability to move from a large unit of measure to a smaller unit of measure simply by moving the decimal point or changing the exponent. 2 -7 For example, 0.003 km is easily changed to 3.00m and 4.50 x 10 nm is easily changed to 4.50 x 10 m by applying a few simple rules. Step 1: Determine the number of decimal places between the units involved in the conversion. * Memorize the chart at the end of this document including prefixes! The most common units are shown in the graph below. You can use the mnemonic King Henry Died by Drinking Chocolate Milk (Kilo, Hecto, Deka, Base, Deci, Centi, Milli) to help remember these. Decimal places from 3 0 -2 -3 -6 -9 base Unit km = 103 m m = 100 m cm = 10-2 m mm = 10-3 m µm = 10-6 m nm = 10-9 Move decimal to the left to convert to a larger unit. Move decimal to the right to convert to a smaller unit Step 2: for Standard Numbers: If you are converting from a large unit to a smaller unit the number will get bigger and the decimal place will move to the right. If you are converting from a smaller unit to a larger unit the number will get smaller and the decimal place will be moved to the left. A way to remember the direction of the decimal shift is to use this mnemonic: Large Unit → Small Unit → Large Number Small Unit → Large Unit → Small Number Example: Convert 0.003km to cm. Step 1: There are 5 decimals between km and cm. (3-(-2)) = 5 Step 2: km is larger than cm so the number must become larger. The decimal must be moved to the right by a total of 5 decimal places. Therefore 0.003km = 300cm Scientific Notation: If you are converting from a large unit to a smaller unit the number becomes larger which means the exponent must increase. If you are converting from a smaller unit to a larger unit the number will become smaller and the exponent will decrease. An easy way to remember the direction of the decimal shift is to use the previously stated rule of thumb:  -3 Example: Convert 3.0 x 10 µm to cm. Step 1: There are 4 decimals between µm and cm. (-6(-2)) = -4) Step 2: µm is smaller than cm which means the number must become smaller! The exponent must -3 -7 be decreased by 4. Therefore 3.00 x 10 µm = 3.00 x 10 cm You can also always use dimensional analysis/factor labeling to do these metric conversions! *********************************************************************************************************************** Section 2: Using Consistent Units in Calculations: When performing calculations, it is important to verify that all of the basic units of measurement (length, mass, time, etc) are measured in the same metric prefix. Example: An ant was observed to travel 3.00m south, turn to the west and move an additional 50.1cm, and finally turn to the north and travel an additional 0.0110km. Determine the total distance in meters traveled by the ant. Solution: The first step is to recognize that the three distances have been given to you in different units of length. Before you can perform the addition you will need to convert all of the measurements to the same unit of length. In this case the most convenient choice is the meter. Make certain to preserve the correct number of significant figures as you make the conversions. 3.00m = 3.00m (3 sf) 50.1cm = 0.501m (3 sf) 0.0110km=11.0 m (3 sf) We can now proceed with the addition: (3.00m + 0.501m + 11.0m) = 14.501m Next use the addition -1 rule (least precise decimal) and round to 10 : Reported Answer: 14.5m ****************************************************************************************** Problems Part (a): Make the following conversions – preserve the number of significant figures in the answer! 1. 450nm ____________________ mm 2. 34km ______________________cm 6 3. 43 000mm _____________________m 4. 4.0 x 10 nm __________________µm -3 5. 3.98 x 10 km ___________________m 6. 456mm ______________________km 12 7. 136 000m _____________________km 8. 4.89 x 10 mm ________________km 6 9. 2.68 x 10 m __________________km 10. 456 000 µm _________________mm 11. 450mm ____________________m 12. 23cm _____________________mm 4 13. 234 µm __________________cm 14. 2.34 x 10 cm _________________m -7 15. 4.56 x 10 cm __________________nm  Unit Multiplication – Dimensional Analysis – Factor Labeling Units: In the world of mathematics numbers often exist as abstract and unit-less entities. However, in the world of physics and chemistry where numbers are based upon experimentation and measurement all numbers are based in a physical reality. As a result, every number consists of two important parts. The first is a magnitude and the second equally important part is a unit. It is the unit that gives physical, real-world meaning to the number. We never write one without the other! Examples: Note that these are all “equivalence statements”! 12 inches in one foot 365 days in one year 7 days in one week 9 1.0 x 10 bytes in one gigabyte Derived Units and Calculations Many of the common units we use are actually derived units that result from performing mathematical operations on the basic units. When performing mathematical operations the units are treated and manipulated as if they were algebraic variables. Here are a few examples: 2 Area = (length -m) x (width -m) = m Volume = (length -m) x (width -m) x (height - m) = m3 Velocity = ( distance traveled – m)/(time -s) = m/s Density = (mass – g)/(volume - mL ) = g/mL Unit Conversions It is often necessary to convert from one system of units to another. The most efficient way to do this is using a process known as “unit multiplication”, “factor labeling” or “dimensional analysis”. Example No. 1: Consider a pin measuring 2.85 cm in length in the metric system. What would be the corresponding length in the English system? Step 1: find an equivalence statement: i.e. 1 inch = 2.54 cm Step 2: Now divide both sides by 2.54 cm: → 1 inch/2.54 cm = 1 or 2.54 cm/1 inch = 1 This gives rise to two conversion factors: Step 3: Chose the conversion factor that will result in the cancellation of the original unit 1 𝑖𝑛𝑐ℎ 2.85 𝑐𝑚 × = 1.12 𝑖𝑛𝑐ℎ𝑒𝑠 2.54 𝑐𝑚 Note that the units for cm cancels out (cm is in both the numerator and denominator) leaving the desired units of inches! “goal posting” One useful version of this method is called “goal posting”. Step 1: Draw a “goal post ”with the horizontal bar extending on each side. Step 2: Place the original number and unit to the left. Place the final unit on the right. Step 3: Move the original unit (cm) from the top left (numerator) to the bottom of the conversion factor (denominator). Now there is no confusion about which form of the conversion factor you will use. If you have done this correctly the original units on the top (cm) will be cancelled by the same unit in the denominator of the conversion factor.  Dimensional Analysis 1. I have 470 milligrams of table salt, which is the chemical compound NaCl. How many liters of NaCl solution can I make if I want the solution to be 0.90% NaCl? (9 grams of salt per 1000 grams of solution). The density of the NaCl solution is 1.0 g solution/mL solution. 2. I have a bar of gold that is 7.0 in  4.0 in  3.0 in. The density of gold is 19.3 g/cm3. The price of gold currently is $1,945.94 per ounce. How much is my gold bar worth? 3. If the RDA for vitamin C is 60 mG per day and there are 70 mg of vitamin C per 100 G of orange, how many 3 oz. oranges would you have to eat each week to meet this requirement? 4. Owls generally maintain territories of 3 acres. How many owls could live in a large wooded area of 20 hectares? (1 hectare=1 sq. dekameter=100 m2= 2.47 acres) 5. The speed of light is 3.00 x 108 m/s. Convert this speed into feet per year. 6. Many candy bars have 9 G of fat per bar. If during a “chocolate attack” you ate one pack of candy (0.6 dekabars), how many ounces of fat would you have eaten? B) There are approximately 9 Calories per gram of fat, how many Calories is this? C) A Calorie is 4184 joules (J). It takes 4.184 J to heat 1 gram of water by 1C. If you wanted to raise the temperature of water by 10C, how many liters of water could you heat with the energy from a pack of candy bars? (Density of water = 1 g/mL) – This one is hard! 7. I have 14.25 ng of glucose (C6H12O6). If 180.18 grams is the mass of 6.10 x 1023 molecules of glucose, how many carbon atoms are in my sample? Part IIIa: Subatomic Particles, Isotopes and Ions Average Element or Atomic Neutrons* (for most common Abbreviation Atomic Mass Protons* Electrons* Ion Number (Z) isotope unless otherwise noted) (A) Oxygen O 8 16.00 Bismuth Bi 209.0 F- Carbon C 6 12.01 14 Carbon-14 C 14.00 6 Pb-208 15 30.97 15 55.845 23 Potassium Ion K+ 39.10 18 (cation) Sulfur Ion S2- 32.07 (anion) *- Calculate the number of protons, neutrons, and electrons for the most prevalent isotope Average Atomic Masses: Silver has two isotopes, one with 60 neutrons and the other with 62 Potassium has three isotopes. The number of neutrons and the neutrons. Give the chemical notation for each of these isotopes and natural abundance of these are: 20 neutron (93.23%); 21 neutrons calculate the relative abundance for each isotope given that the average (0.012%); and 22 neutrons (6.73%). Give the chemical notation for atomic mass for silver is 107.87 amu. each of these isotopes and calculate the average atomic mass for potassium.  Part IIIb: Electron Configuration & Orbital Diagrams In the space below, write the electron configurations of the following elements: 1. Oxygen ________________________________________________ 2. Chlorine ________________________________________________ 3. Sodium ________________________________________________ 4. Aluminum ________________________________________________ 5. Argon ________________________________________________ 6. Iron ________________________________________________ 7. Potassium ________________________________________________ 8. Scandium ________________________________________________ 9. Bromine ________________________________________________ 10. Barium ________________________________________________ 11. Iodine ________________________________________________ 12. Strontium ________________________________________________ 13. Yttrium ________________________________________________ 14. Cadmium ________________________________________________ 15. Tin ________________________________________________ Determine what elements are denoted by the following electron configurations:  Part IIIb: Electron Configuration & Orbital Diagrams 11) 1s22s22p63s23p5 ____________________ 12) 1s22s22p63s23p64s23d104p65s24d1 ____________________ 13) 1s22s22p3 ____________________ 14) 1s22s22p63s23p64s23d104p65s24d105p66s24f145d6 ____________________ 15) 1s22s22p63s23p64s23d104p3 ____________________ Draw the orbital diagrams for the flowing elements: Example: Mg (12 e-) __ __ __ __ __ __ 1s 2s 2p 3s 16) Nitrogen 17) Sodium 18) Chlorine 19) Potassium 20) Iron 21) Zinc 22) Selenium 23) Ruthenium 24) Antimony 25) Xenon Chemistry I Name: _____________________________ Block: ______ Part IV: Periodic Trends 1. On the blank periodic table, color and label: a. alkali metals b. alkaline metals c. transition metals d. nonmetals e. metalloids f. halogens g. noble gases h. inner transition metals 2. On the blank periodic table, color and label. a. the s block b. the p block c. the d block f. the f block 3. On the blank periodic table, draw arrows to show the following periodic trends across each period and down each group. Be sure to label which way the trend is increasing and which way it is decreasing. a. Atomic radius b. Ionization energy c. Electronegativity  1 18 1 Periodic Table of the Elements 2 13 14 15 16 17 2 3 3 4 5 6 7 8 9 10 11 12 4 5 6 * 7 ** * ** ©2015 Todd Helmenstine sciencenotes.org  Part IV: Periodic Trends Worksheet Directions: Use your notes to answer the following questions. 1. Rank the following elements by increasing atomic radius: carbon, aluminum, oxygen, potassium. 2. Rank the following elements by increasing electronegativity: sulfur, oxygen, neon, aluminum. 3. Why does fluorine have a higher ionization energy than iodine? 4. Why do elements in the same family generally have similar properties? 5. Indicate whether the following properties increase or decrease from left to right across the periodic table. a. atomic radius (excluding noble gases) b. first ionization energy c. electronegativity 6. What trend in atomic radius occurs down a group on the periodic table? What causes this trend? 7. What trend in ionization energy occurs across a period on the periodic table? What causes this trend? 8. Circle the atom in each pair that has the largest atomic radius. a. Al or B c. Na or Al e. S or O b. O or F d. Br or Cl f. Mg or Ca 9. Circle the atom in each pair that has the greater ionization energy. a. Li or Be c. Ca or Ba e. Na or K b. P or Ar d. Cl or Si f. Li or K 10. Define electronegativity. 11. Circle the atom in each pair that has the greater electronegativity. a. Ca or Ga c. Br or As e. Li or O b. Ba or Sr d. Cl or S c. O or S  Part V: Chemical Bonding Section 1: Ionic Bonding Ionic bonds involve a transfer of electrons from one atom (or atomic group) to another. Cations are positive ions resulting from the loss of electrons. Anions are negative ions resulting from the gain of electrons. Atoms generally lose or gain electrons to achieve a “stable octet” or set of 8 electrons in the valence shell (although there are exceptions!) Metals tend to have low electronegativity and ionization energy and tend to form cations. Nonmetals tend to have high electronegativity and tend to form anions. Things to know – study the charts available on the course website! 1. Placement of metals and nonmetals on Periodic Table. 2. The charges/oxidation states taken by elements in different groups of Periodic Table. 3. Charges of common metals that take multiple charges (multivalent metals). 4. Common Polyatomic Ions (memorize the chart – both names and formulas with charges!). Section 2: Covalent Bonding Covalent bonds involve a sharing of electrons between atoms. Usually both elements in a covalent bond are nonmetals. Equal sharing of electrons produces a nonpolar covalent bond and occurs when the bonding atoms have equal or very similar electronegativity. Unequal sharing of electrons occurs when atoms have significantly different electronegativities and results in a polar covalent bond in which one atom has a partial negative charge and the other a partial positive charge. Things to know: 1. Be able to determine whether a bond is ionic, polar covalent or nonpolar covalent based on the elements bonding and electronegativity chart. 2. Draw a basic Lewis Dot structure showing the placement of all electrons. Bonding occurs on a spectrum based on the difference in electronegativity between the two atoms involved in the bond. Memorize the rules below and have a general sense of the electronegativities of common elements (& how the trend runs along the periodic table)! Difference in electronegativity 0 0.5 1.0 2.0 4.0 Moderately Polar Very Polar-covalent Nonpolar Covalent Ionic bond Covalent bond Rules of thumb: ∆EN > 2.0 → Bond is ionic ∆EN < 0.5 → Bond is nonpolar covalent 0.5 ≤ ∆EN ≤ 1. 6 → Bond is polar covalent 1.6  ∆EN ≤ 2.0 → Bond is polar covalent IF it involves two nonmetals, otherwise ionic. Sharon Karackattu Problems! More Less Difference in Bond Bonding between electronegative electronegative electronegativity Type element and value element and value Sulfur & Hydrogen Sulfur and cesium Chlorine and bromine Calcium and chlorine Oxygen and hydrogen Nitrogen & hydrogen Iodine and iodine Copper and Sulfur Hydrogen & Fluorine Carbon and Oxygen Sharon Karackattu  Part VI: Nomenclature of Binary Compounds ** Before you start naming compounds or writing formulas from names be sure to review which elements are metals, transition metals & nonmetals and the charges they take as well as common polyatomic ions with their charges (makes this much easier!) Part 1: Determine if the compound is ionic or covalent to decide which set of naming rules to apply: A. Ionic compound: i. Compound contains a polyatomic ion ii. Compound contains a metal and a nonmetal B. Covalent compound: i. Compound contains only nonmetal elements Part 2: Ionic Compound Nomenclature A. Name the cation i. Univalent metal cations = same name as the element a. Na+ = sodium, Ba2+ = barium, Al3+ = aluminium etc. b. These are usually Group 1, 2 and 13 elements ii. Multivalent metal cations = same name as element + charge denoted by Roman Numeral in parenthesis a. Fe2+ = Iron (II), Fe3+ = Iron (III) b. Multivalent metal cation are usually in the transition metal block (Iron, Copper, Nickel, Chromium etc.) c. Silver is always 1+ (Ag+) so it has no Roman Numeral d. Zinc is always 2+ (Zn2+) so it has no Roman Numeral e. An easy way to remember charges for Al, Zn and Ag is noting that they form a diagonal step down starting with Al going down to the left (3+, 2+ and 1+) f. Pb and Sn are two metals not in the transition block that can take either the charge 2+ or 4+. As such, Pb and Sn always have a Roman Numeral when being named in a compound. iii. If the cation is a polyatomic ion – it takes the same name as the ion. I.e. NH4+ is ammonium. B. Name the anion i. Anion that is based on a nonmetal element: a. Use the root of the elemental name b. Change the suffix to -ide c. Cl- = chloride, O2- = oxide, P3- = phosphide, N3- = nitride etc. ii. Anion that is a polyatomic ion: a. Use the name of the polyatomic ion b. SO42- = sulfate, PO33- = phosphite, CrO42- = chromate etc. Sharon Karackattu  C. Examples: MgCl2 = magnesium chlorid FeCl3 = iron (III) chloride NH4Cl = ammonium chloride Sn3(PO4)2 = Tin (II) phosphate (NH4)2SO4 = ammonium sulfate Part 3: Covalent Compound Nomenclature A. Name the first element – use Greek Prefixes (except mono) i. Select the appropriate Greek prefix using subscript of the element a. Mono = one b. Di = two c. Tri = three d. Tetra = four e. Penta = five f. Hexa = six g. Hepta = seven h. Octa = eight i. Nona = nine j. Deca = ten ii. Name the first element using the prefix and the element name: a. Do not use the prefix mono- for the first element. If there is only one atom of the first element in the compound “mono” is implied B. Name the second element i. Select the appropriate Greek prefix using the subscript of the element ii. Use the root of the element name for the second element iii. Convert the suffix of the elemental name to -ide. C. Examples: H2O = dihydrogen monoxide (the o from mono- gets dropped in monoxide) CO2 = carbon dioxide CO = carbon monoxide PCl5 = phosphorus pentachloride S2O3 = disulfur trioxide Sharon Karackattu  Naming Binary Metals or Polyatomic Ions Involved? Chemical Compounds Yes No Ionic Covalent Multivalent Cation Monovalent Cation Polyatomic Cation 1. Name cation by element 1. Name cation by 1. Name cation by name element name name of polyatomic → Use Roman Numeral in 1. 1st Greek prefix (don’t Ex. Na+ = Sodium, cation parentheses to denote use mono) Ca2+= Calcium, Ex. Ammonium charge 2. Name first element Ag+ = Silver Ex. Fe2+ = Iron (II), Fe3+= Iron (III) 3. 2nd Greek prefix 4. Root of 2nd element Monoatomic Anion Polyatomic Anion 5. Change suffix to –ide 2. Name anion by 2. Name anion by element root. polyatomic anion Examples: 3. Change suffix to -ide name carbon monoxide dinitrogen tetraoxide Examples: Cation + Monoatomic Anion phosphorus pentachloride sodium fluoride, calcium bromide, ammonium chloride, iron (II) oxide sulfur hexafluoride Examples: Cation + Polyatomic Anion dihydrogen monoxide sodium phosphate, ammonium carbonate, copper (II) sulfate dihydrogen dioxide Names to Formulas of Metals or Polyatomic Ions Involved? Chemical Compounds Yes No Ionic Example – iron (III) sulfate Covalent 1. Use the name to determine the two ions in the compound → Fe and SO42- 2. Write the cation first (remember Roman Numeral = charge on metal cation). 1. 1st Greek prefix denotes Then write the anion. Include charges (for now) → Fe3+SO42- subscript of first element 2. Write element symbol and 3. Balance the charges on the two ions to obtain a neutral formula unit. The easy subscript way is to “criss-cross” so that the charge on the cation becomes the subscript of the anion. The charge of the anion becomes the subscript on the cation. Use the lowest whole number ratio of subscripts! → Fe3+2SO42-3 3. 2nd Greek prefix denotes subscript of second element 4. If the subscript of a polyatomic ion is greater than 1, put the whole polyatomic 4. Write symbol and subscript ion symbol in parentheses and the subscript outside the parenthesis. → for second element Fe3+2(SO42-)3 5. Erase any ion charges in the formula → Fe2(SO4)3 Examples: carbon monoxide = CO Examples: Cation + Monoatomic Anion dinitrogen tetraoxide = N2O4 sodium fluoride = NaF , calcium bromide = CaBr2, sulfur hexafluoride = SF6 ammonium chloride = AlCl3, iron (II) oxide = FeO, iron (III) oxide = Fe2O3 dihydrogen monoxide = H2O Examples: Cation + Polyatomic Anion dihydrogen dioxide = H2O2 Copper (ii) phosphate = Cu3(PO4)2 , ammonium carbonate = (NH4)2CO3 carbon tetrahydride = CH4 Part VI: Problems - More Naming Practice! Ionic or Covalent? vanadium (V) phosphate ____________________________ __________________ sodium permanganate ______________________________ __________________ MnF2 ____________________________________________ __________________ Ni(SO3)2 _________________________________________ __________________ phosphorus triiodide _______________________________ __________________ H3PO4 ___________________________________________ __________________ HI ______________________________________________ __________________ Pb3N4 ___________________________________________ __________________ Sn(OH)2 ________________________________________ __________________ SiCl4 ____________________________________________ __________________ HClO2 ___________________________________________ __________________ Sodium sulfate ____________________________________ __________________ Hydrosulfuric acid __________________________________ __________________ Nitrogen trifluoride __________________________________ __________________ Calcium phosphide _________________________________ __________________ B2Si _____________________________________________ __________________ PCl5 _____________________________________________ __________________ Perbromic acid ______________________________________ __________________ Manganese (IV) carbonate ____________________________ __________________ C2H4 ______________________________________________ __________________ Carbon disulfide _____________________________________ __________________ Iron (III) nitrate ______________________________________ __________________ Copper (II) phosphite _________________________________ __________________ Sulfur hexachloride ___________________________________ __________________  Write the Name or the Chemical Formula Antimony tribromide ________________ Aluminum sulfide ____________________ Lithium oxide _____________________ P4S5 ______________________________ Tin (II) hydroxide ___________________ chlorine dioxide _____________________ B2Si ____________________________ NF3 ______________________________ Iron (III) phosphide _________________ Cobalt (III) carbonate _________________ Hydrogen iodide ____________________ SeF6 ______________________________ Zn3(PO4)2 ________________________ Be(NO3)2 __________________________ Dinitrogen trioxide __________________ Na2(SO3)3 _________________________ Sodium hydroxide __________________ Iodine pentafluoride ___________________ Cu(CH3 COO)2 ____________________ Hexaboron silicide ____________________ Si2Br6 ____________________________ Cu(HCO3)2 _________________________ Phosphorus triiodide _________________ CH4 _______________________________ Writing Chemical Formulas Practice I Fill in the symbols and charges of the ions and then write the correct chemical formulas and the chemical names in the corresponding blocks. The first one is done for you. IONS Sodium Calcium Aluminum Ammonium Hydrogen Na + Chloride NaCl Cl - Sodium chloride Acetate Oxide Sulfite Phosphate Iodide  Part VII: Mole Conversions Notes & Practice Worksheet There are three mole equalities. They are: 1 mol = 6.02 x 1023 particles 1 mol = molar mass in grams (periodic table) 1 mol = 22.4 L for a gas at STP Each equality can be written as a set of two conversion factors. They are: 𝟏 𝒎𝒐𝒍𝒆 𝟔.𝟎𝟐 × 𝟏𝟎𝟐𝟑 𝒑𝒂𝒓𝒕𝒊𝒄𝒍𝒆𝒔 𝟏 𝒎𝒐𝒍𝒆 𝒎𝒐𝒍𝒂𝒓 𝒎𝒂𝒔𝒔 𝒊𝒏 𝒈𝒓𝒂𝒎𝒔 ( ) 𝒐𝒓 ( ) ( ) 𝒐𝒓 ( ) 𝟔.𝟎𝟐 × 𝟏𝟎𝟐𝟑 𝒑𝒂𝒓𝒕𝒊𝒄𝒍𝒆𝒔 𝟏 𝒎𝒐𝒍𝒆 𝒎𝒐𝒍𝒂𝒓 𝒎𝒂𝒔𝒔 𝒊𝒏 𝒈𝒓𝒂𝒎𝒔 𝟏 𝒎𝒐𝒍𝒆  22.4 L    or  1 mole    at Standard Temperature and Pressure (0C and 1 atm)  1 mole   22.4 L  Example Problems: 1. How many moles of magnesium is 3.01 x 1022 atoms of magnesium?  1 mole  3.01 x 1022 atoms   = 5 x 10-2 moles  6.02 x10 atoms  23 2. How many molecules are there in 4.00 moles of glucose, C6H12O6?  6.02 x1023 molecules  4.0 oles   = 2.41 x 1024 molecules  1 mole  3. How many moles in 28 grams of CO2 ? Molar mass of CO2 1 C = 1 x 12.01 g = 12.01 g 2 O = 2 x 16.00 g = 32.00 g 44.00 g/mol  1 mole  28 g CO2   = 0.64 moles CO2  44.00 g  4. What is the mass of 5 moles of Fe2O3 ? Molar mass Fe2O3 2 Fe = 2 x 55.6 g = 111.2 g 3 O = 3 x 16.0 g = 48.0 g 5.2 g/mol  159.2 g  5 moles Fe2O3   = 800 grams Fe2O  1 mole  5. Determine the volume, in liters, occupied by 0.030 moles of a gas at STP.  22.4 L  0.030 mol   = 0.67 L  1 mole   6. How many moles of argon atoms are present in 11.2 L of argon gas at STP?  1 mole  11.2 L   = 0.500 moles  22.4 L  Mixed Mole Conversion Examples: Given unit → Moles → Desired unit 7. How many oxygen molecules are in 3.36 L of oxygen gas at STP?  1 mole   6.02 x1023 molecules  3.36 L     = 9.03 x 1022 molecules  22.4 L  1 mole  8. Find the mass in grams of 2.00 x 1023 molecules of F2 Molar mass 2 F = 2 x 19 g = 38 g/mol  1 mole   38 g  2.00 x 1023 molecules  23    = 12.6 g  6.02 x10 particles   1 mole  ************************************************************************************************************* Problems I: Mole Conversions Practice – Show Work 1. How many moles are 1.20 x 1025 atoms of phosphorous? 2. How many atoms are in 0.750 moles of zinc? 3. How many molecules are in 0.400 moles of N2O5? 4. Find the number of moles of argon in 452 g of argon. 5. Find the grams in 1.26 x 10-4 mol of HC2H3O2. 6. Find the mass in 2.6 mol of lithium bromide. 7. What is the volume of 0.05 mol of neon gas at STP? 8. What is the volume of 1.2 moles of water vapor at STP? 9. Determine the volume in liters occupied by 14 g of nitrogen gas at STP. 10. Find the mass, in grams, of 1.00 x 1023 molecules of N2. 11. How many particles are there in 1.43 g of a molecular compound with a gram molecular mass of 233 g? 12. Aspartame is an artificial sweetener that is 160 times sweeter than sucrose (table sugar) when dissolved in water. It is marketed by G.D. Searle as Nutra Sweet. The molecular formula of aspartame is C14H18N2O5. a) Calculate the gram molar mass of aspartame. b) How many moles of molecules are in 10 g of aspartame? c) What is the mass in grams of 1.56 moles of aspartame? d) How many molecules are in 5 mg of aspartame? e) How many atoms of nitrogen are in 1.2 grams of aspartame?  Chemical Reactions Review Sheet Types of Chemical Reactions: Combination or Synthesis A + B → AB Decomposition AB → A + B Single Replacement A + BC → B + AC Double Replacement AB + CD → AD + CB Can be a) acid-base if the reactants are acid & base and products are salt & water. b) can be precipitation if a solid product forms Hydrocarbon Combustion CxHyOz + O2 → CO2 + H2O Oxidation-Reduction - Involve a transfer of electrons. Occurs during combustion, single replacement and can occur during synthesis and decomposition. Problems: 1. A reaction occurs when aqueous lead (II) nitrate is mixed with an aqueous solution of potassium hydroxide. Write an overall, balanced equation for the reaction, including state designations. 2. For the following three reactions, label the type, predict the products (make sure formulas are correct), and balance the equation. __________________ a. _____C3H4 (g) + _____O2 (g) → __________________ b. _____ Ba(NO3)2(aq) + ______ Na3PO4(aq) → __________________c. _____ Al(s) + ______ O2(g) → __________________d. ______ HBr(aq) + ______ KOH(aq) → __________________e. ______ Ca(NO3)2 (aq) + ______ Na2SO4 (aq) → 3. In the following equations, label the oxidized element and the reduced element. a. 2Na(s) + Cl2(g) → 2NaCl(s) b. 2NaBr(aq) + Cl2(g) → 2NaCl(s) + Br2(l)  Reaction Quest Review 1. What are 4 signs that a reaction is taking place? Think back to the lab: 2. What is does it mean when a substance is reduced? When it is oxidized? How is a single replacement reaction an oxidation-reduction reaction? 3. What are the 5 main types of chemical reactions? What type of reaction is an acid-base neutralization? 4. What does (s), (g), (l) and (aq) mean when placed near a chemical formula in an equation? A) WRITE THE FORMULA FOR EACH MATERIAL CORRECTLY. B) BALANCE THE EQUATION. SOME REACTIONS REQUIRE COMPLETION. C) FOR EACH REACTION TELL WHAT TYPE OF REACTION IT IS. D) For double and single replacement reactions – write the net ionic equations. 1. sulfur trioxide and water combine to make sulfuric acid. 2. lead II nitrate and sodium iodide react to make lead iodide and sodium nitrate. 3. calcium fluoride and sulfuric acid (H2SO4) make calcium sulfate and hydrofluoric acid 4. calcium carbonate decomposes when you heat it to leave calcium oxide and carbon dioxide. 5. ammonia gas when it is pressurized into water will make ammonium hydroxide. 6. sodium hydroxide neutralizes carbonic acid 7. zinc sulfide and oxygen become zinc oxide and sulfur. 8. lithium oxide and water make lithium hydroxide 9. aluminum hydroxide and sulfuric acid neutralize to make water and aluminum sulfate. 10. sulfur burns in oxygen to make sulfur dioxide. 11. barium hydroxide and sulfuric acid make water and barium sulfate. 12. aluminum sulfate and calcium hydroxide become aluminum hydroxide and calcium sulfate. 13. copper metal and silver nitrate react to form silver metal and copper II nitrate. 14. propane burns (with oxygen) 15. zinc and copper II sulfate yield zinc sulfate and copper metal 19. sulfuric acid reacts with zinc 22. calcium oxide and aluminum make aluminum oxide and calcium  Name__________________________________ Period_____ Net Ionic Equation Worksheet READ THIS: When two solutions of ionic compounds are mixed, a solid may form. This type of reaction is called a precipitation reaction, and the solid produced in the reaction is known as the precipitate. You can predict whether a precipitate will form using a list of solubility rules such as those found in the table below. When a combination of ions is described as insoluble, a precipitate forms. There are three types of equations that are commonly written to describe a precipitation reaction. The molecular equation shows each of the substances in the reaction as compounds with physical states written next to the chemical formulas. The complete ionic equation shows each of the aqueous compounds as separate ions. Insoluble substances are not separated and these have the symbol (s) written next to them. Water is also not separated and it has a (l) written next to it. Notice that there are ions that are present on both sides of the reaction arrow –> that is, they do not react. These ions are known as spectator ions and they are eliminated from complete ionic equation by crossing them out. The remaining equation is known as the net ionic equation. For example: The reaction of potassium chloride and lead II nitrate Molecular Equation: 2KCl (aq) + Pb(NO3)2 (aq) -> 2KNO3 (aq) + PbCl2 (s) Complete Ionic Equation: 2K+ (aq) + 2Cl- (aq) + Pb2+ (aq) + 2NO3– (aq) -> 2K+ (aq) + 2NO3– (aq) + PbCl2 (s) Net Ionic Equation: 2Cl- (aq) + Pb2+ (aq) -> PbCl2 (s) Directions: Write balanced molecular, ionic, and net ionic equations for each of the following reactions. Assume all reactions occur in aqueous solution. Include states of matter in your balanced equation. 1. Sodium chloride and lead II nitrate Molecular Equation: Net Ionic Equation: 2. Sodium carbonate and Iron II chloride Molecular Equation: Net Ionic Equation: 3. Ammonium phosphate and zinc nitrate Molecular Equation:  Net Ionic Equation: 4. Iron III chloride and magnesium metal Molecular Equation: Net Ionic Equation: 5. Silver nitrate and magnesium iodide Molecular Equation: Net Ionic Equation: 6. Aluminum and copper (II) perchlorate Molecular Equation: Net Ionic Equation: 7. Sodium and water Molecular Equation: Net Ionic Equation: 8. Zinc and hydrochloric acid Molecular Equation: Net Ionic Equation:  Steps to Find Empirical & Molecular Formulas Remember this: “Percent to mass, Mass to mole, Divide by small, Make it whole” 1. Determine the mass in grams of each element present in the sample. “Percent to mass” If the information in the problem is in terms of percent composition of each element → a) assume you have 100 g of the sample to start with b) The grams of each element (out of the 100 g sample) will just be the numerical value of its percent composition. EXAMPLE: You have a sample that is 40.0% carbon, 6.73% hydrogen and the rest oxygen. Find the empirical and molecular formulas. Step 1: 40.0% + 6.73% = 46.73%. The percentage of oxygen is 100%-46.73% = 53.27% If I have 100 g of sample to start with, I have: 40.0 grams Carbon, 6.73 grams Hydrogen and 53.27 grams Oxygen 2. Calculate the number of moles of each element. “Mass to mole” Step 2: Moles of Carbon = 40.0g C x 1 mol C/12.01g C = 3.331 mol C Moles Hydrogen = 6.73g H x 1 mol H/1.01g = 6.663 mol H Mole Oxygen = 53.27 g O x 1 mol O/16.0 g = 3.33 mol O DO NOT ROUND THESE NUMBERS → KEEP SEVERAL DECIMAL PLACES 3. Divide each by the smallest number of moles to obtain the simplest whole number ratio. “Divide by small” Step 3: The molar ratio of the elements in my compound is C3.331H6.663O3.33. I want a whole number ratio, so I will divide all the subscripts by the smallest number of moles (3.331) to get: C1H2O1 → so my empirical formula is CH2O If your number after dividing are values like 2.07, 1.1 etc. then round to the nearest whole number. If they are values like 3.5, 2.333 etc., then go to step 4. Sharon Karackattu 4. If whole numbers are not obtained* in step 3), multiply through by the smallest integer that will give all whole numbers “Make it whole” Let’s say that my empirical formula turned out to be C2.333H4O2. 2.333 is not close enough to 2 to round down to 2. But I can multiply my formula through by 3 to get this: C7H12O6 5. Finding molecular formula: If the molar mass of your empirical formula matches the molar mass of the final compound (as stated in the problem) → Hooray! You are done: your empirical formula IS your molecular formula. Step 5: For my example in step 1, it says that the molecular weight (molar mass) of my compound is 180.18 g/mol My empirical formula is CH2O from step 3 has a molar mass of (12.01 + 2×1.01 + 16) g/mol = 30.03 g/mol. So my empirical formula is not my molecular formula. Now, divide molar mass of compound/molar mass of empirical formula: 180.18 g/mol ÷ 30.03 g/mol = 6 The molar mass of my compound is 6 times the molar mass of my empirical formula. Multiply the empirical formula subscripts by 6 to get the final molecular formula: 6(CH2O) = C6H12O6 → The compound in my sample is glucose! Sharon Karackattu  Steps to Solving Limiting Reagent Problems Suppose 13.7 g of C2H2 reacts with 18.5 g O2 according to the reaction below. What is the mass of CO2 produced? What is the limiting reagent? 2C2H2(g) + 5O2(g) → 4CO2(g) + 2H2O(ℓ) 1. Find the mass of product yielded by the given amount of the first reactant. You can use either product (CO2 or H2O), but since the question asks about CO2, it will be easier to use this product: 13.7 g C2H2 1 mole C2H2 4 mole CO2 44.02 g CO2 = 46.3 g CO 26.04 g C2H2 2 mole C2H2 1 mole CO2 2. Find the mass of the same product (in this case CO2) yielded by the given amount of the second reactant. 18.5 g O2 1 mole O2 4 mole CO2 44.02 g CO2 = 20.4 g CO2 32.00 g O2 5 mole O2 1 mole CO2 3. Since the 18.5 grams of O2 produces less CO2, it is the limiting reagent in this problem. This amount of O2 gets used up first and “limits” how much CO2 can be produced. The amount of CO2 that can be produced is 20.4 grams (which you already calculated!) 4. You can repeat steps 1 and 2 for any number of reactants that you have a given mass for. The limiting reagent will ALWAYS be the reactant that produces the least amount of product (because it gets used up first). 5. Finding the amount of excess reagent: The excess reagent is the one that is NOT the limiting reagent. There will be some of this reagent leftover after the limiting reagent is completely used up. Figure out how much of the excess reagent must react completely with the given amount of the limiting reagent. Then subtract this amount from the given amount of the excess reagent. 18.5 g O2 1 mole O2 2 mole C2H2 26.02 g C2H2 = 6.02 g C2H2 used 32.00 g O2 5 mole O2 1 mole C2H2 13.7 g of C2H2 total – 6.02 g of C2H2 used = 7.68 g C2H2 excess (leftover) Sharon Karackattu  Part VIII: Stoichiometry-Based Problems 1. a) Nicotine is a stimulant and an addictive chemical found in tobacco. An analysis of nicotine produces the following percent composition: 74.03% carbon, 17.27% nitrogen, and 8.70% hydrogen. What is the empirical formula of nicotine? b) Further tests show that the molar mass of nicotine is 162.23 g/mol. Given this information, what is the molecular formula of nicotine? 2. An ionic sample with a mass of 0.5000 g is determined to contain the elements indium and chlorine. If the sample has 0.2404 g of chlorine, what is the empirical formula of this ionic compound? 3. A 16.4 g sample of hydrated calcium sulfate is heated until all the water is driven off. The calcium sulfate that remains has a mass of 13.0 g. Find the formula and the chemical name of the hydrate. 4. ____C3H8 + _____O2 → a. What type of reaction is written above? _________________ b. Predict the products of the reaction and balance it. c. If I start with 5.00 grams of C3H8 and 5.00 grams of O2, what is the limiting reagent? What is my theoretical yield of the carbon containing product?  d. I get a percent yield of 75%. How many grams of the carbon containing product did I make? 5. Magnesium undergoes a single replacement reaction with hydrochloric acid. a) Write the Balanced Equation: b) Which element is oxidized? _______ Which element is reduced? ______ c) How many grams of hydrogen gas can be produced from the reaction of 3.00 g of magnesium with 4.00 g of hydrochloric acid? d) Identify the limiting and excess reactants. How many grams of the excess reagent are leftover? e) If the hydrogen gas is produced at 48C and 2.5 atm of pressure, what is the volume produced in liters? 6. Sulfur reacts with oxygen to produce sulfur trioxide gas. a) Write the Balanced Equation: b) If 6.3 g of sulfur reacts with 10.0 g of oxygen, what is the theoretical yield of sulfur trioxide gas in grams? c) What is the limiting reagent? How many grams of the excess reagent is leftover? d) The sulfur trioxide gas produced had a volume of 5.4 L and was produced at 98C. What is the pressure of the gas in kPa?  Part IX: Gas Laws, Molarity, pH and Putting it all Together 1. The following questions pertain to the reaction below: ______HBr + _____ Ca → a. What type of reaction is shown above? ____________________ b. Predict products and then balance the reaction. c. Name the ionic product of the reaction. ____________________ d. Which element is oxidized? _________ Which element is reduced? ________ e. 1.7 grams of Ca are mixed with 850.6 mL of 0.043 M HBr. What is the maximum theoretical yield of the gaseous product in grams? f. How many grams of the excess reagent are leftover? g. What is the pH of the HBr solution? h. What is the OH- concentration of the HBr solution? i. If the gas is produced at 89C and 1.7 atm of pressure, what is the volume of gaseous product in mL? j. The pressure of the gas is changed to 250 mmHg and the volume is changed to 1.54 L. What is the temperature of the gas now? Question 2: The following questions pertain to the reaction below ____H3PO4(aq) + _____ Ca(OH)2(aq) → a) What type of reaction is shown above? _______________________________ (HINT: It could be two of the types we learned about because one product is insoluble – which one? ___________). b) Predict the products and balance the reaction. c) Write the net ionic reaction for the reaction above. d) Name the reactants and products. Identify acid, base, conjugate acid and conjugate base. e) If I have 7.62 grams of Ca(OH)2, what volume of 0.050 M H3PO4 would be required to react with it completely? f) In the reaction, only 6.89 grams of the solid product were produced. What is the percent yield of the reaction? g) How many grams of the Ca(OH)2 remained unreacted? Question 3: It takes combustion of 58.8 mL of liquid propane (C3H8), which has a density of 0.493 g/cm3, to cook my hamburger. If air is 21.0% by volume O2, how many liters of air at 27.0C and 105.0 kPa will it take to cook my burger? (NOTE: this is not happening at STP!) a) Write and balance the combustion reaction for propane b) Calculate the grams of propane used to cook the burger c) Calculate the moles of oxygen used to cook the burger d) Calculate the volume of O2 used to cook the burger e) Calculate the volume of air used to cook the burger