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AP Topic 1 Review 1. Classify the following as an element, a compound or a mixture, if you classify it as a mixture state whether it is a homogeneous or heterogeneous mixture. i) Helium ii) Pure water iii) Pure table salt (sodium chloride) iv) Air v) Fruit cake 2. Classify the following as either chemical or physical changes. i) Ice melting ii) Gasoline burning 3. Mercury is a liquid metal that has a density of 13.58 g/mL. Calculate the volume of mercury that must be poured out in order to obtain 0.5g of mercury. 4. Classify the following as either quantitative or qualitative observations. i) My eyes are brown ii) My neck size is 17 inches iii) My average grade last year was 79% 7. Round the following numbers to four figures. i) 2.16347 x 105 ii) 4.000574 x 106 iii) 3.682417 iv) 7.2518 v) 375.6523 vi) 21.860051 D:\769818713.doc Page 1 of 17 8. Perform the following conversions. i) 0.75 kg to milligrams ii) 1500 millimeters to km 9. Complete the following table of average temperatures on the planets. Kelvin Mercury 439 Venus 729 Fahrenheit Celsius Earth 15 AP Topic 2 Review 1. Using the periodic table, complete the following table. (13) Isotope Symbol Atomic # Mass # 38 88 # Protons # Neutrons 35 44 23Na 11 2. What is the charge on a sodium atom? __________ 3. What is the charge on a sodium nucleus? __________ 4. What is the atomic number of potassium? __________ 5. How many protons in the nucleus of a potassium atom? __________ 6. How many electrons in the potassium nucleus? __________ 7. Define the term isotope. D:\769818713.doc Page 2 of 17 # Electrons 8. Using the periodic table, a) Complete the following table. Isotope Symbol Atomic # # Protons # Neutrons Mass # 13C 17 18 26 56 17 37 2 52 3 128 50 70 Naturally occurring lead is found to have the following isotopic relative abundance. 204Pb 3%, 206Pb 24%, 207Pb 20% and 208Pb 53%. Calculate the average relative atomic mass of Pb from the data. The results taken from a mass spectrum of chlorine gas show peaks at m/e 35.00 and m/e 37.00 (The mass spectrum identifies the different isotopes of an element that are present in a sample). a) Given that the relative abundances of Cl 35.00 and Cl 37.00 are 77.50% and 22.50% respectively, calculate the average relative atomic mass of chlorine to four significant figures. Give the name of each of the following binary compounds. a) KCl b) BaO c) Rb2S d) Na3P e) AlF3 f) Mg3N2 g) CaI2 h) RaCl2 D:\769818713.doc Page 3 of 17 Write the name of the following ionic substances using the system that includes a Roman numeral to specify the charge on the cation. a) FeI3 b) MnCl2 c) HgO d) Cu2O e) CuO f) SnBr4 g) MnO2 Write the name of each of the following binary compounds of non-metals. a) N2Br4 b) P2S5 c) SeO2 d) N2O5 e) SiO2 Write the name of each of the following acids. a) HCl b) HNO2 c) H2SO4 d) HF e) HI f) HC2H3O2 AP Topic 3 Review 1. Give full and abbreviated (noble gas core) electronic configurations for the following. a) Cr b) Fe c) S2- D:\769818713.doc FULL ______________________________ NOBLE GAS CORE ______________________________ FULL ______________________________ NOBLE GAS CORE ______________________________ FULL ______________________________ NOBLE GAS CORE ______________________________ Page 4 of 17 2. For each of the following sets of orbitals, indicate which orbital is higher in energy. a) 1s, 2s _________ b) 2p, 3p _________ c) 4s, 3dyz _________ d) 3px, 3py, 3pz _________ 3. Indicate the block (s, p or d) in which each of the following elements found. BLOCK a) Sc ______ b) P ______ c) Fr ______ d) Ni ______ e) As ______ 4. Identify the element from the electron configurations of atoms shown below. a) [Ne] 3s2 3p2 _________ b) [Ar] 4s2 3d7 _________ c) [Xe] 6s2 _________ 5. State which atom or ion is represented by the following sets of atomic numbers and electronic configurations. Atomic # Electronic Configuration a) 8 1s2 2s2 2p4 ________ b) 11 1s2 2s2 2p6 ________ c) 14 1s2 2s2 2p6 3s2 3p2 ________ d) 22 1s2 2s2 2p6 3s2 3p6 3d2 ________ 6. Give the electron configurations for the following transition metal ions. a) Sc3+ __________________________________ b) Cr2+ __________________________________ c) Ni3+ __________________________________ D:\769818713.doc Page 5 of 17 7. Identify two positive and two negative ions that are isoelectronic with argon. Two Positive ions ________ ________ Two Negative ions ________ ________ 8. Using the electrons in boxes notation complete the electronic configurations of the following elements. 1s 2s 2p 3s 3p 3d 4s 4p Element V 1s 3s 4s 2s 2p 3p 3d 4p Element Ar 1s 3s 4s 2s 2p 3p 3d 4p Element Zn 9. State the number of unpaired electrons in each of the electronic configurations in question 9 # of unpaired electrons V ________ Ar ________ Zn ________ 10. Write three possible sets of quantum numbers for the highest energy electrons in the aluminum atom. n l ml ms Electron # 11 Electron # 12 Electron # 13 11. Which atomic theory is violated by the following sets of quantum numbers representing beryllium’s outer shell electrons? Explain your answer. n 2 2 L 0 0 ml 0 0 ms +½ +½ _____________________________________________________________________________ _____________________________________________________________________________ D:\769818713.doc Page 6 of 17 AP Topic 4 Review TYPE 1: Those involving Avogadro’s number. Question 1 A sample of Ag is found to contain 9.7 x 1023 atoms Ag. How many moles of Ag atoms are in the sample? Question 2 How many Sb atoms are found in 0.43 moles of pure Sb? TYPE 2: Those involving the relationship between mass, moles and molar mass. Question 3 What is the mass in grams of 2.53 moles Al? Question 4 How many moles of Na in 20g of Na? Question 5 If 50 moles of a simple, binary, group I chloride have a mass of 3725g identify the group I metal. TYPE 3: Those combining types 1 & 2. Question 6 How many Zr atoms are found in a 1.23g sample of Zr? Question 7 What is the mass of 5.14 x 1023 atoms of uranium? Question 8 What mass of C atoms have the same number of atoms as are in a 11.2g sample of Si? D:\769818713.doc Page 7 of 17 TYPE 4: % by mass Composition. Question 9 Calculate the percent by mass composition of ethanol, C2H6O. Question 10 What is the percent by mass composition of N2O5? Question 11 A compound has the formula Al4[Fe(CN)6]3. What is the percent by mass composition of this compound? TYPE 5: Empirical formulae. Question 12 A compound contains 26.9% N and 73.1% F. What is the empirical formula of the compound? Question 13 2.3g of magnesium is completely reacted with 6.75g of chlorine. What is the empirical formula of the compound formed? TYPE 6: Molecular formulae from empirical formulae. Question 14 What is its molecular formula of hydrocarbon that has an empirical formula of C 2H5 and a molecular mass of 58. Question 15 A compound contains 68.54% carbon, 8.63% hydrogen, and the remainder oxygen. The molecular weight of this compound is approximately 140g/mol. What is the empirical formula? What is the molecular formula? D:\769818713.doc Page 8 of 17 Type 7: Combustion Analysis. Question 16 The combustion of 2.95 grams of a compound that contains only C, H and S yields 5.48 grams of CO2 and 1.13 grams of H2O. What is the empirical formula of the compound? Type 8: % Yield. In questions 17 assume that the reactant that data is supplied for, is limiting. Question 17 If 31 grams of C4H10 produces 41 grams of CO2 what is the % yield? Type 9: Limiting Reactant. Question 18 Consider the reaction between Aluminum and Iron (III) oxide to produce Aluminum oxide and Iron metal. a) Write an equation for the reaction. b) If 1240g of Al are reacted with 6010g of Iron (III) oxide, identify the limiting reagent. Which reagent is in excess? c) Calculate the mass of Iron formed. D:\769818713.doc Page 9 of 17 d) How much of the excess reagent is left over at the end of the reaction? Type 10: Analysis of hydrated salts. Question 19 Copper (II) sulfate is found as a hydrated salt, CuSO 4.xH2O. A technician carefully heats 2.50g of the salt to a constant mass of 1.60g. a) What is meant by constant mass? b) How many moles of copper sulfate are there in 1.60g of anhydrous copper (II) sulfate? c) How many moles of water were lost? d) What is the value of x in the formula? Type 11: Moles and reacting ratios (including solutions). Question 20 Sodium hydrogen carbonate, NaHCO3, combines with HCl as indicated below. NaHCO3(aq) + HCl(aq) NaCl(aq) + CO2(g) + H2O(l) a) What volume of 1.5 M HCl solution should be present to combine totally with 0.14 moles of NaHCO3? b) How many moles of CO2 are produced when 0.49 g of NaHCO3 combines with excess HCl? D:\769818713.doc Page 10 of 17 c) Calculate the mass of NaCl that results when 1.48 moles of HCl combines with excess NaHCO3. d) What mass of NaHCO3 is required to produce 6.1 x 103 moles of H2O? Type 12: Dilution. Question 24 Calculate the volume of 2M sulfuric acid that must be diluted with water to produce 2dm3 of 0.5M sulfuric acid. AP Topic 5 Review In all questions show all relevant working and balance any equations you write. 1. Electrolytes & non-electrolytes Indicate if you would expect the following compounds to be electrolytes or non-electrolytes when in aqueous solution. In each case very briefly explain your answer. Use equations if appropriate. (6) a) Sodium chloride b) Methanol (CH3OH, an alcohol similar to ethanol – NOT a hydroxide) c) Magnesium nitrate D:\769818713.doc Page 11 of 17 3. Precipitation reactions and net ionic equations Predict if a precipitation reaction will occur in each of the following cases. If it does, write the full, balanced equation AND the net ionic equation (including state symbols) to show the formation of the precipitate. If there is no reaction, say so, and indicate why. (9) a) CuSO4(aq) + Na2CO3(aq) b) HI(aq) + Zn(NO3)2(aq) c) AgNO3(aq) + NaBr(aq) 4. Acids and bases & neutralization a) Write a full, balanced equation for the reaction of aqueous sulfuric acid and aqueous potassium hydroxide to produce an aqueous salt and water. b) Rewrite the equation in b) removing any spectator ions (i.e. write the net ionic equation) with state symbols. c) If I used 34.89 mL of a 0.345 M solution of sulfuric acid to neutralize a 20.00 mL sample of potassium hydroxide, what is the molarity of the KOH solution? 5. Oxidation number concept What is the oxidation number of each of the underlined atoms in each of the following species? Think carefully about the rules that are being applied and write a very brief, simple explanation of your answer in each case. a) CaI2 b) GeO2 c) KO2 d) NH3 6. REDOX - Balance the following equations Zn + D:\769818713.doc HCl Zn2+ + H2 Page 12 of 17 (acidic) Cr2O7 + Cl- MnO4- + S2- Cr3+ + MnS + Cl2 S (acidic) (basic) 8. REDOX Titration calculations a) Consider the following half equations. SO32- + H2O SO42- + 2H+ + 2eMnO4- + 8H+ + 5e- Mn2+ + 4H2O Combine these two equations to obtain the overall reaction of sulfite ions with manganate (VII) ions. b) Use the equation you have written in a) to calculate the volume of 0.222 M manganate (VII) ions that are required to react completely with 25.0 cm 3 of 0.456 M sulfite ions. 9. Classification of chemical reactions By choosing two of the following reaction types from the list below, classify each of the following reactions in two ways. Reaction types: precipitation, acid-base, REDOX (oxidation and reduction), single displacement, double displacement, combination, decomposition, combustion. a) 2NaOH(aq) + CuSO4(aq) Na2SO4(aq) + Cu(OH)2(s) AND b) HI(aq) + KOH(aq) KI(aq) + H2O(l) AND c) C3H8(g) + 5O2(g) 3CO2(g) + 4H2O(l) AND d) 2HgO(s) 2Hg(l) + O2(g) D:\769818713.doc Page 13 of 17 AND e) 4B(s) + 3O2(g) 2B2O3(s) AND f) 2Al(s) + 6HNO3(aq) 2Al(NO3)3(aq) + 3H2(g) AND AP Topic 6 Review 1. Small quantities of hydrogen gas can be prepared in the laboratory by the following reaction: Zn(s) + 2HCl(aq) ZnCl2(aq) + H2(g) Assume you carried out this experiment and collected 653 mL of hydrogen gas over water. The gas mixture collected includes hydrogen and water vapor. The temperature of this gas mixture was 23.0 °C and the total pressure was 892 mm Hg. How many moles of hydrogen did you prepare? The vapor pressure of water at 23.0°C is 19.8 mm Hg. 2. A mixture of H2(g), O2(g), and 2 milliliters of H2O(l) is present in 0.500-liter rigid container at 25 °C. The number of moles of H2 and the number of moles of O2 are equal. The total pressure is 1,146 millimeters of mercury. (The equilibrium vapor pressure of pure water is 24 millimeters mercury.) The mixture is sparked, and H2 and O2 react until one reactant is completely consumed. a. Identify the reactant remaining and calculate the number of moles of the reactant remaining. b. Calculate the total pressure in the container at the conclusion of the reaction if the final temperature is 90 °C. (The equilibrium vapor pressure of water at 90 °C is 526 millimeters mercury.) D:\769818713.doc Page 14 of 17 c. Calculate the number of moles of water present as vapor in the container at 90 °C. AP TOPIC 7 Review 2. Define ionization energy. (2) _____________________________________________________________________________ _____________________________________________________________________________ 3. Using the metal magnesium, write an equation to summarize the process of first ionization energy. (Remember state symbols are important as they from part of the definition). (2) _____________________________________________________________________________ 4. Which of the following pairs has the largest size? (3) PAIR 1 Ca, Ca2+ ____________________________________ PAIR 2 Na, Mg ____________________________________ PAIR 3 S, S2- ____________________________________ 5. Put the following species into order of increasing ionization energy (smallest first). (2) Ca, Ca+, Ca2+, Sr _____________________________________________________________________________ 6. Given that the first ionization energy of sodium is 496 kJmol -1. Calculate (showing your working and logic) the energy required to turn 2.3g of sodium from gaseous atoms to gaseous +1 ions. (2) 7. Explain carefully why rubidium tends only to form a +1 ion? (2) _____________________________________________________________________________ _____________________________________________________________________________ 8. Explain carefully why elements in the same group react in similar ways? (1) _____________________________________________________________________________ _____________________________________________________________________________ D:\769818713.doc Page 15 of 17 9. Explain carefully the factor, when moving up and down groups I & II, that determines the pattern of reactivity that is observed? (2) _____________________________________________________________________________ _____________________________________________________________________________ 10. Write equations to illustrate the following. (i) The basic nature of potassium oxide. (2) _____________________________________________________________________________ (ii) The acidic nature of both the common oxides of sulfur. (4) (SO2)_________________________________________________________________________ (SO3)_________________________________________________________________________ (iii) The amphoteric nature of aluminum oxide. (4) (ACIDIC CHARACTER)__________________________________________________________ (BASIC CHARACTER)__________________________________________________________ 11. Using the following data estimate the boiling point of bromine. (1) ELEMENT Boiling Point/K Fluorine Chlorine Iodine 85 239 458 _____________________________________________________________________________ 12. Define electron affinity. (2) _____________________________________________________________________________ _____________________________________________________________________________ 13. Write an equation to summarize the process of second electron affinity of oxygen. (2) _____________________________________________________________________________ 14. Consider the table of ionization energies for element X shown below. Ionization Energy in kJ/mol 1st 2nd 3rd 4th 5th 6th 737 1450 7732 10540 13360 17995 (i) In which group will X be found? (1) ___________________________________________ (ii) Explain your answer to (i). (2) _____________________________________________________________________________ _____________________________________________________________________________ D:\769818713.doc Page 16 of 17 (iii) Predict the formula of X’s bromide. (1) __________________________________________ Topic 8 Review Complete the Table AROUND CENTRAL ATOM Molecule or ion # of Bonding Areas # of Lone pairs Atom geometry SF6 PCl5 SO3-2 CO32- D:\769818713.doc Page 17 of 17 Lewis dot structure and hybridization