AP Chemistry Summer Review Winston Churchill High School

Transcription

1 Summer Review Dear AP CHEM STUDENTS, We are looking forward to the school year and the work we re going to do together. This packet is meant to refresh you on several things you already learned so that you have them firmly in mind for the start of school. We recommend that you review your Honors Chemistry notes, and start looking at this material several weeks before school starts in the fall. For turning the assignment in you should copy and paste the questions into a separate file which you then print, and from there you can complete the questions. If you have any questions during the summer regarding the assignment Mr. Nugent (Mrs. Bopanna will be enjoying herself in beautiful India for the summer) will be available at Churchill from the second week of July through the first week of August, and otherwise available at the following addresses: Chemistry_frisson@verizon.net Or James_M_Nugent@mcpsmd.org This Assignment will be collected and graded after you return to school in the fall. Good luck, and have a wonderful and productive summer! Sincerely, J. Boppana, J. Nugent 1

2 Summer Review Chem Sheets to Memorize SOLUBILITY CHART Soluble Exceptions Insoluble Exceptions No NO 3 Smoking S 2 Group IA, NH + 4, Sr 2+, Ba 2+, Ca 2+ CHeating CH 3 COO Counter 2 CO 3 Group IA, + NH 4 Cellphones Cl Ag +, Pb 2+ 3, Productive PO 4 Hg2 2+ Group IA, NH 4 + Bullying Br Ag +, Pb 2+, 2+ Hg 2 Intimidating I Ag +, Pb 2+, 2+ Hg 2 O (ZERO) OH Group IA, NH + 4, Sr 2+, Ba 2+, Ca 2+ Students SO 4 2 Sr 2+, Ba 2+, Ca 2+, Pb 2+, Hg 2 2+ Always Soluble Generally Soluble Exceptions Generally Insoluble Exceptions Group IA ions Cl, Br, I Ag +, Pb 2+, Hg2 2+ S 2, OH Group IA, NH + 4, Sr 2+, Ba 2+, Ca 2+ NH 4 + NO 3 CH 3 COO F Ca 2+, Sr 2+, Ba 2+, Pb 2+, Mg 2+ CO 3, PO 3 4, SO 2 3, C 2 O 2 4, CrO 2 4, O 2 Sr 2+, Ba 2+, Ca 2+, Pb 2+ 2+, Hg 2 SO 4 2 Group IA, NH 4 + ClO 3, ClO 4 2

3 Summer Review Acids HF weak HCl strong HBr strong HI strong H 2 SO 4 strong HNO 3 strong HClO 3 strong HClO 4 strong All else are weak H 2 CO 3 H 2 O + CO 2 (very weak acidbreaks down!) Bases LiOH strong NaOH strong KOH strong (all IA Metal Hydroxides) Ca(OH) 2 strong Sr (OH) 2 strong Ba(OH) 2 strong Special Reactions Metal metal + acid salt + H 2 metal oxide + H 2 O metal hydroxide metal oxide + CO 2 metal carbonate metal chloride + O 2 metal chlorate Nonmetal nonmetal oxide + H 2 O oxyacid SO 3 + H 2 O H 2 SO 4 SO 2 + H 2 O H 2 SO 3 N 2 O 5 + H 2 O 2 HNO 3 N 2 O 3 + H 2 O 2 HNO 2 P 2 O 5 + 3H 2 O 2 H 3 PO 4 P 2 O 3 + 3H 2 O 2 H 3 PO 3 3

4 Summer Review Oxidizers Acid Base Neutral MnO 4 or MnO 2 Mn 2+ Cr 2 O CrO 4 MnO 4 MnO 2 2 CrO 4 Cr 3+ Cr 2 O 2 7 Cr 3+ NO 3 (dil) NO NO 3 (conc) NO 2 metallic ions metallous ions free halogens halide ions Na 2 O 2 NaOH HClO 4 Cl C 2 O 2 4 CO 2 H 2 O 2 O 2 + H 2 O Reducers halide ions free halogens free metals metal ions metallous ions metallic ions SO 3 2 SO 4 2 NO 2 NO 3 free halogens (dil) hypohalite ions free halogens (conc) halate ions Colors of Complex Ions 2+ [Cr(H 2 O) 6] [Cu(H 2 O) 6 ] 2+ [Cu(H 2 O) 6 ] 3+ [Mn(H 2 O) 6 ] 2+ [Co(H 2 O) 6 ] 2+ [Fe(H 2 O) 6 ] 2+ [Ni(H 2 O) 6 ] 2+ [Fe(H 2 O) 6 ] 3+ Ion Blue Blue Blueviolet Very pale pink Pink Pale green Green Yellowbrown Color 4

5 Summer Review Flame Test Colors Li +, Sr 2+, Ca 2+ Na + K + Ba 2+ Cu 2+ Ion Color Red (various shades) Yellow/Orange Lilac Green Bluegreen Transition Metal Ion Colors Sc Oxidation Number colorless Ti VIOLET colorless V VIOLET GREEN BLUE YELLOW Cr BLUE GREEN YELLOW(CrO 4 2 ) ORANGE(Cr 2 O 7 2 ) Mn PALE PINK BROWN DARK GREEN DARK PURPLE Fe PALE GREEN REDDISH BROWN Co PINK YELLOW ORANGE Ni GREEN Cu colorless BLUE Zn colorless 5

6 Summer Review Common Precipitate Colors WHITE BLUE YELLOW BLACK GREEN REDDISH BROWN AgCl Many Copper (II) precipitates AgI Many Sulfides Many Iron(II) precipitates Many Fe (III) precipitates BaSO 4 PbI 2 PbCl 2 Many nontransition metal hydroxides Many nontransition metal carbonates and sulfates Common Tests for Gases GAS Hydrogen gas Oxygen gas Carbon dioxide gas Ammonia gas TEST Squeaky pop with lighted splint Relights glowing splint Turns limewater (Calcium Hydroxide solution) milky Pungent odor, turns red litmus paper blue, gives dense white fumes in contact with conc. HCl fumes 6

7 Summer Review Common Tests for Cations and Anions ION Carbonate and Hydrogen carbonate Sulfate Chloride Bromide Iodide Ammonium TEST Release CO 2 gas with acids White ppt of BaSO 4 with barium ions White of AgCl with silver ions Cream ppt of AgBr with silver ions Yellow ppt of AgI with silver ions NH 3 released with hydroxide ions Halogen Colors Fluorine (gas) Chlorine (gas) Bromine (liquid) Iodine (solid) NO 2 (gas) Pale yellowgreen Green Orangebrown Dark purple Orangebrown Color Changes in Redox reactions 1) MnO 4 (aq) Mn 2+ (aq) (Dark Purple) (Pale Pink) 2) CrO 4 2 (aq) Cr 3+ (aq) (Orange) (Green) Acid/Base Indicator Color Changes INDICATOR ACID BASE Methyl Orange RED YELLOW Methyl Red RED YELLOW Litmus RED BLUE Universal RED BLUE/PURPLE 7

8 Summer Review Phenolphthalein COLORLESS PINK Polyatomic Ions Group IIIB or 13 Group IVB or 14 Group VB or 15 Charge 3 Charge 2 Charge 1 NO 3 3 BO 3 2 CO 3 Nitrate Borate Carbonate NO 2 Nitrite one member in the ion family Remember: ions with the greater # of oxygens: ATE ions with the fewer # of oxygens: ITE adding hydrogen in front makes BI and reduces charge by 1 SiO 3 2 Silicate PO 4 3 Phosphate PO 3 3 Phosphite AsO 4 3 Arsenate AsO 3 3 Arsenite two members in the ion family Charge 3 Other Important Polyatomic Ions to Remember Group VIB or 16 SO 4 2 Sulfate SO 3 2 Sulfite O SeO 4 2 Selenate SeO 3 2 Selenite Group VIIB or 17 F ClO 4 Perchlorate ClO 3 Chlorate ClO 2 Chlorite ClO Hypochlorite BrO 4 Perbromate BrO 3 Bromate BrO 2 Bromite BrO Hypobromite IO 4 2 TeO 4 Periodate Telurate IO 3 2 TeO 3 Iodate Telurite IO 2 Iodite IO Hypoiodite Charge 2 Charge 1 four members in the ion family Acetate C 2 H 3 O 2 Chromate 2 CrO 4 Hydroxide OH Dichromate 2 Cr 2 O 7 Permanganate MnO 4 Peroxide 2 O 2 Cyanide CN Oxalate 2 C 2 O 4 Hydronium H 3 O + Thiosulfate 2 S 2 O 3 Ammonium + NH 4 Tartrate 2 C 4 H 4 O 6 Bisulfite HSO 3 Bisulgate HSO 4 Bicarbonate HCO 3 Biphosphite 2 HPO 3 Biphosphate 2 HPO 4 Hydrogen H 2 PO 3 8

9 Summer Review Thiocyanate SCN biphosphite Copy and paste the questions on pages 9 through 18 and answer the following questions on that copy. Your responses do not need to be typed, but you should write your answers neatly and legibly. Nomenclature: Naming And Writing Formulas Of Chemical Compounds Formula Name 1. P 4 O ZnAt 2 3. SBr 6 4. CaF 2 5. P 2 S 3 6. carbon monoxide 7. sodium hydride 8. aluminum selenide 9. xenon hexafluoride 10. dinitrogen monoxide 11. KClO Pb(OH) Ca(MnO 4 ) N 2 O Ti(HPO 4 ) manganese (VII) oxide 17. francium dichromate 18. copper (II) dihydrogen phosphate 19. silver chromate 20. ammonium oxalate 21. (NH 4 ) 2 SO Ni 3 (PO 4 ) Fe(IO 2 ) NaBrO H 3 PO tartaric acid 27. hydrotelluric acid 9

10 Summer Review 28. mercury (I) nitrate Significant Figures 1. Give the number of significant figures in each of the following numbers\ a. 123 b c d. 12,000 e. 1,000,000,000.0 f g. 23,000. h. 34,000 i j Do the following calculations giving the answer in the appropriate number of sig figs a b c * 8.2 d. 234 / e f g * 3 h. 25,600 / Do the following calculations giving the answer in the appropriate number of sig figs a * / 75 b. ( ) * (23, ) c * / 0.1 d. (8+9) / ( ) Stoichiometry & General Chemistry You will need to be able to write molecular chemical reactions and do mole conversions to do the following questions: g of sodium metal reacts with a solution of excess lithium bromide. How many grams of lithium metal are produced? 2. How many molecules are in 100.L of 0.1M potassium hydroxide solution at STP? 3. Propane, C 3 H 6, undergoes combustion. How many grams of propane are needed to produce 45.9g of water? 4. How many moles or water are in 3.02 X molecules of water? 5. Find the empirical and molecular formulas for a compound containing g iron and 5.01 g oxygen if the molar mass of the compound is 320g/mol 6. A solution of 3.50g of sodium phosphate is mixed with a solution containing 6.40g of barium nitrate. How many grams of barium phosphate can be formed? 7. Find the empirical and molecular formulas for a compound containing 5.28g of tin and 3.37g of fluorine if the molar mass of the compound is 584.1g/mol 10

11 Summer Review 8. Octane, C 8 H 18, undergoes combustion. How many grams of oxygen are needed to burn 10.0g of octane? 9. Sodium azide, NaN 3, decomposes into its elements. How many grams of sodium azide are required to form 34.6g of nitrogen gas. 10. Ammonia reacts with oxygen gas to form nitrogen monoxide and water. How many grams of nitrogen monoxide are formed when 1.50 g of ammonia react with 2.75 g of oxygen gas? Reactions Write balanced net ionic equations (NIEs), with states of matter included, i.e. (g), (s), (aq), and (l), for the following questions on a separate piece of paper. You ll have reactions that are classified as precipitation, acidbase, or redox (reductionoxidation, which includes synthesis, decomposition, and single displacement/replacement). Notes: Any ion has an aqueous state of matter. For acidbase reactions, strong acids (HCl, HBr, HI, H 2 SO 4, HClO 4, and HNO 3 ) and strong bases (metal ions in groups 1 and 2 paired with hydroxide (OH ) completely dissociate. Weak acids and bases do not. For precipitation (and some redox) reactions, use the solubility rules below to determine which salts are soluble (aqueous) or insoluble (solid). Only aqueous solutions can dissociate solids, liquids, and gases cannot. 1. Salts containing Group I elements are soluble (Li +, Na +, K +, Cs +, Rb + ). Exceptions to this rule are rare. Salts containing the ammonium ion (NH 4 + ) are also soluble. 2. Salts containing nitrate ion (NO 3 ) and acetate ion (C 2 H 3 O 2 ) are generally soluble. 3. Salts containing Cl, Br, F, and I are generally soluble. Important exceptions to this rule are halide salts of Ag +, Pb 2+, and Hg Thus, AgCl, PbBr 2, and Hg 2 Cl 2 are all insoluble. 4. Most silver salts are insoluble. AgNO 3 and AgC 2 H 3 O 2 are common soluble salts of silver; virtually anything else is insoluble. 5. Most sulfate salts are soluble. Important exceptions to this rule include BaSO 4, PbSO 4, Ag 2 SO 4 and SrSO Most hydroxide salts are only slightly soluble. Hydroxide salts of Group I elements are soluble. Hydroxide salts of Group II elements (Ca, Sr, and Ba) are slightly soluble. Hydroxide salts of transition metals and Al 3+ are insoluble. Thus, Fe(OH) 3, Al(OH) 3, and Co(OH) 2 are not soluble. 7. Most sulfides of transition metals are highly insoluble. Thus, CdS, FeS, ZnS, and Ag 2 S are all insoluble. Arsenic, antimony, bismuth, and lead sulfides are also insoluble. 8. Most chromates, phosphates, bicarbonates, and carbonates are frequently insoluble except those with alkali metals and ammonium. 11

12 Summer Review AcidBase Example: Hydrochloric acid is added to a solution containing zinc hydroxide. This problem is a little bit trickier than your standard acidbase neutralization problem. That said, it s worth paying attention to how we look at this problem and what state symbols we use, specifically with Zn(OH) 2, for the reaction between the HCl and the Zn(OH) 2. First, write an equation representing the reaction: HCl + Zn(OH) 2 ZnCl 2 + H 2 O Next, determine what dissociates and what does not. Hydrochloric acid is a strong acid, so it will completely dissociate into its ions, while zinc hydroxide is a weak base (all group one metals and group two metals strontium and below are considered strong bases, everything else is a weak base), so it will not dissociate on its own in solution. The reaction that occurs with zinc hydroxide in water is as follows: Zn(OH) 2 Zn OH This is an equilibrium reaction, not a complete dissociation of the hydroxide as you will have with HCl which dissociates nearly 100% in water. There will be a slight dissociation for the zinc hydroxide, but not anywhere near 100%. That said if you recall Le Chatlier s principle: If the equilibrium of a system is disturbed by a change in one or more of the determining factors (as temperature, pressure, or concentration) the system tends to adjust itself to a new equilibrium by counteracting as far as possible the effect of the change. In this example we will disturb the equilibrium by removing OH (per the definition above, the determining factor that we re changing is the concentration of OH in solution) as we add HCl to the solution as the H + ion will react with the OH, creating H 2 O in the process. As we remove the OH in the reaction with the acid, we force the reaction above to move to the right, creating more OH in the process. In effect we force the Zn(OH)2 to become soluble by adding acid to the base and neutralizing whatever little hydroxide actually goes into solution with water. Zinc chloride is a soluble salt according to the solubility rules above, so it will also dissociate into its ions. Wait to balance the reaction until the end. 12

13 Summer Review H + + Cl + Zn(OH) 2 Zn 2+ + Cl + H 2 O Last, you need to see what can be cancelled out. Species that are identical on both sides of the reaction, called spectator ions, can be cancelled out. Cl is present on both sides of the reaction and therefore can be cancelled out giving you your net ionic reaction that you ll now balance and put back on states of matter. 2 H + (aq) + Zn(OH) 2 (s) Zn 2+ (aq) + 2 H 2 O (l) Redox Example: Zinc metal reacts with a solution of silver nitrate. 2AgNO 3 (aq) + Zn (s) Zn(NO 3 ) 2 (aq) + 2Ag (s) The potential difference between zinc and silver is such that zinc will be induced to send its electrons to silver, turning zinc into an ion and silver into a solid (the details of how this works will be dealt with later in the course.) Breaking this down into a net ionic equation (often referred to as an NIE): Zn + 2Ag + + 2NO 3 2Ag + Zn NO 3 *NO 3 is a spectator ion and we therefore remove it from the net ionic equation as it doesn t change in the course of the reaction: Zn + 2Ag + 2Ag + Zn 2+ Precipitation Example: Barium acetate is mixed with potassium sulfate. Ba(C 2 H 3 O 2 ) 2 + K 2 SO 4 BaSO 4 + KC 2 H 3 O 2 *According to the solubility rules, barium sulfate is the only insoluble salt. So, everything else will dissociate. Ba 2+ + C 2 H 3 O 2 + K + + SO 4 2 BaSO 4 + K + + C 2 H 3 O 2 *The potassium ions and acetate ions can be cancelled out. Ba 2+ (aq) + SO 4 2 (aq) BaSO 4 (s) Complete the following questions on a separate piece of paper: 1. Solid sodium bicarbonate is mixed with copper (II) nitrate 2. Magnesium oxide is heated 13

14 Summer Review 3. Acetic acid is added to a solution of ammonia 4. Iron (III) chloride is mixed with silver sulfite 5. A solid piece of aluminum is put into a solution of nickel (II) chloride 6. A solution of lithium chloride is added to a solution of lead (IV) nitrite 7. Sulfuric acid is added to a solution of aluminum hydroxide 8. Cadmium nitrate is added to sodium sulfide 9. Chromium (III) sulfate is added to ammonium carbonate 10. Methane combusts in air In each of the equations below, the reactants are written correctly. You must write the correct products and then balance the equation. Identify the type of chemical reaction before writing the products. 1. CaCO3 2. Al + O2 3. Fe + CuSO4 4. C6H12 + O2 5. Zn + H2SO4 6. Cl2 + MgI2 7. NaOH 8. Fe + HCl 9. NaOH + H3PO4 10. (NH4)2SO4 + Ca(OH)2 11. AgNO3 + K2SO4 12. Mg(OH)2 + H3PO4 13. Na + H2O 14. KClO3 15. Al2(SO4)3 + Ca3(PO4)2 16. SO2 + H2O 17. (NH4)3PO4 + Ba(OH)2 14

15 18. Ca(OH)2 + HNO3 19. C3H8 + O2 20. Li + S AP Chemistry Summer Review Electron Structure and Periodicity You will need to know about valence electrons, electron shells, orbital notation, electron configuration, atomic radius, ionization energy, and electronegativity to do these questions. 1. Draw the orbital notation for nickel. 2. How many unpaired electrons are in Arsenic? 3. Write the electron configuration for palladium 4. How many valence electrons are in mercury? 5. Write the electron configuration for uranium 6. Write the noble gas electron configuration for lead. 7. Which is more electronegative, sulfur or chlorine? Why? 8. Which has a larger atomic radius, potassium or bromine, and why? 9. Which has the smaller ionization energy, nitrogen or phosphorous, and why? 10. Write the noble gas electron configuration for copper. Short Answer questions from previous AP Exams 11. Use the principles of atomic structure and/or chemical bonding to explain each of the following. In each part, our answer must include references to both substances. a) The atomic radius of Li is larger than that of Be b) The second ionization energy of K is greater than the second ionization energy of Ca c) The carbontocarbon bond energy in C 2 H 4 is greater than it is in C 2 H 6 d) The boiling point of Cl 2 is lower than the boiling point of Br 2 15

16 Summer Review Atomic Structure Sample problems 12. Give the symbols for the isotopes of carbon, nitrogen and uranium and determine the number of protons, electrons, and neutrons in each isotope. 13. Given the data below determine the average atomic mass of the elements shown: Isotope % Abundance Isotopic Mass a. Sb % amu Sb % amu b. Ag % amu Ag % amu 14. Convert each of the following to moles Mole Concept Sample Problems a g NaOH b X atoms Au c. 40.0L of Ne gas d. 800.g CaBr 2 e X molecules of H 2 O f. 6.78L of Ar gas 15. Do the following a. Given moles of krypton determine (i) the mass, (ii) the number of atoms and (iii) the volume at STP b. Given moles of oxygen determine (i) the mass, (ii) the number of atoms, and (iii) the volume at STP Bonding *You will need to know about Lewis structures, covalent bonding, shape names, and bond angles to do these questions. *For the following questions, draw the Lewis Structure, name the shape, and state the bond angle. 1. SeCl 2 6. NH CO 2 2. NO 3 7. CO 2 3. OF 2 8. CH 3 NH 2 4. BF 3 9. HCOOH 2 5. SO HCN 16

17 Summer Review Short Answer Problems from Previous AP Exams 1. The reaction between silver ion and solid zinc is represented by the following equation: 2Ag + (aq) + Zn (s) Zn 2+ (aq) + 2Ag (s) A 1.50g sample of Zn is combined with 250mL of 0.110M AgNO 3 at 25 C. a. Identify the limiting reagent. Show calculations to support your answer. b. On the basis of the limiting reactant that you identified in part a, determine the value of [Zn 2+ ] after the reaction is complete. 2. Consider the hydrocarbon pentane, C 5 H 12 (molar mass 72.15g). a. Write the balanced equation for the combustion of pentane to yield carbon dioxide and water. b. What volume of dry carbon dioxide, measured at 25 C and 785mmHg, will result from the complete combustion of 2.50g pentane? 3. Find the mass percent of nitrogen in each of the following compounds: a. NO b. NO 2 c. N 2 O 4 d. N 2 O 4. Benzene contains only carbon and hydrogen and has a molar mass of 78.1g/mol. Analysis shows the compound to be 7.74%H by mass. Find the empirical and molecular formulas of benzene. 5. Calcium carbonate decomposes upon heating, producing calcium oxide and carbon dioxide. a. Write a balanced chemical equation for this reaction. b. How many grams of calcium oxide will be produced after 12.25g of calcium carbonate is completely decomposed? c. What volume of carbon dioxide gas is produced form this amount of calcium carbonate, at STP? 6. Hydrogen gas and bromine gas react to form hydrogen bromide gas. a. Write a balanced chemical equation for this reaction. 17

18 Summer Review b. 3.2g of hydrogen gas and 9.5g of bromine gas react. Which is the limiting reagent? c. How many grams of hydrogen bromide gas can be produced using the amounts in b? d. How many grams of the excess reactant is left unreacted? e. What volume of HBr, measured at STP, is produced in b? 7. When ammonia gas, oxygen gas and methane gas (CH 4 ) are combined, the products are hydrogen cyanide gas and water. a. Write a balanced chemical equation for this reaction. b. Calculate the mass of each product produced when 225g of oxygen gas is reacted with an excess of the other two reactants. c. If the actual yield of the experiment in (b) is 105g of HCN, Calculate the percent yield. 8. When solutions of potassium iodide and lead (II) nitrate are combined, the products are potassium nitrate and lead (II) iodide. a. Write a balanced equation for this reaction, including (aq) and (s). b. Calculate the mass of precipitate produced when 50.0mL of 0.45M potassium iodide solution and 75mL of 0.55M lead (II) nitrate solution are mixed. c. Calculate the volume of 0.50M potassium iodide required to react completely with 50.0mL of 0.50M lead (II) nitrate. 9. Five beakers each containing ml of an aqueous solution are placed on a lab bench. The solutions are all at 25 C. Solution 1 contains 0.20 M KNO 3. Solution 2 contains 0.10 M BaCl 2. Solution 3 contains 0.15 M C 2 H 4 (OH) 2. Solution 4 contains 0.20 M (NH 4 ) 2 SO 4. Solution 5 contains 0.25 M KMnO 4 (for water K f = 1.86 ºC kg/mol and K b = 0.51 ºC kg/mol). a. Calculate the boiling point and freezing point of each solution. b. Explain the cause of the changes in boiling and freezing point. 18