Mr. Reger 1 st Semester Exam Review Date Period. Ch. 2 Measurements and Calculations

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1 Chemistry Name Mr. Reger 1 st Semester Exam Review Date Period Ch. 2 Measurements and Calculations 1. How many significant figures are there in each of the following measurements? a) 6.38 cm b) L c) km d) 5070 g e) 800. J f) 9.30 x 10 4 ml g) x 10-9 m 2. Write the following numbers in scientific notation. a) 360,000 b) c) 9,200,400 d) Write the following as ordinary numbers. a) 6.2 x 10 7 b) x 10-4 c) 3.09 x Write the following numbers with 2 significant figures. a) 18.3 b) c) 146, 400 d) 4, Perform the following conversions. a) 56.8 in = m b) 420. kg = μg c) 28 C = K 6. Perform the following calculations, and write the answer with the correct number of significant figures. a) ml ml ml b) 3.8cm x 2.14cm x 6.31 cm 7. Calculate the density of a piece of metal if its mass is 364g and its volume is 39mL. 1

2 Ch. 3 Matter 1. Identify each of the following as physical properties/changes or chemical properties/changes. a) A piece of gallium metal melts in your hand. b) A piece of wood burns in a fireplace. c) Water is boiled. d) Iron turns into rust. e) A mixture of sand and water is filtered to isolate the sand. f) Solid sodium reacts with chlorine gas to produce solid sodium chloride. 2. Define or describe the following terms: a) Matter b) Compound c) Mixture d) Homogeneous e) Heterogeneous f) Filtration g) Distillation Ch 4. Chemical Foundations: Elements, Atoms, and Ions 1. List the subatomic particles, and give the relative mass, charge, and location (in the atom) of each. 2. Atoms with the same number of protons but different numbers of neutrons are called. 3. How many protons, neutrons, and electrons are there in each of the following? Write the isotope symbol for each. a) a neutral sodium-22 atom b) a Cl - ion with a mass number of 35 c) oxygen-18 d) a neutral atom with an atomic number of 15 and a mass number of Give the name of the group (or section) in the period table to which each element belongs. a) bromine b) potassium c) iron d) krypton e) beryllium 5. Give the element that is located in the following position in the Periodic Table. a) Group 1, period 3 b) Group 4A, period 2 6. List the 4 properties of metals. 2

3 Ch. 5 Nomenclature 1. Name each compound. a) Na 2 O b) Ni(NO 3 ) 2 c) Pb 2 O d) SF 6 e) Mn(OH) 2 f) Cr 2 O 3 g) N 2 Br 4 h) HNO 3 i) PCl 5 j) HBr 2. Write formulas for the following compounds. a) sulfuric acid b) copper(i) oxide c) cadmium chloride d) cesium phosphate e) xenon hexafluoride f) potassium dichromate g) lithium sulfide h) iron(iii) carbonate i) diphosphorous pentoxide j) chromium(ii) fluoride 3

4 Chapter 6 Chemical Reactions: An Introduction 1. Balance the following chemical equations. a) Cl 2 (g) + KI(aq) I 2 (s) + KCl(aq) b) CaF 2 (s) + H 2 SO 4 (aq) CaSO 4 (s) + HF(g) c) C 5 H 12 (g) + O 2 (g) CO 2 (g) + H 2 O(g) d) PbCl 2 (aq) + K 2 SO 4 (aq) PbSO 4 (s) + KCl(aq) e) Fe(NO 3 ) 3 (aq) + KOH(aq) Fe(OH) 3 (s) + KNO 3 (aq) Ch. 7 Reactions in Aqueous Solutions -solubility rules (p.172) -net ionic equations -classifying rxns: A-B, SR, DR, ppt, redox, combustion, decomp, syn 1. Write net ionic equations for the following reactions (predict the products first!). a. BaCl 2 (aq) + Na 2 SO 4 (aq) b. KOH(aq) + HCl(aq) c. Zn(s) + HCl(aq) 3. Classify each rxn in as many ways as possible. a. 4 NH 3 (g) + 5 O 2 (g) 4 NO + 6 H 2 O(g) b. S 8 (s) + 8 O 2 (g) 8 SO 2 (g) c. 2 Al(s) + 3 Cl 2 (g) 2 AlCl 3 (s) d. 2 AlN(s) 2 Al(s) + N 2 (g) e. BaCl 2 (aq) + Na 2 SO 4 (aq) BaSO 4 (s) + 2 NaCl(aq) f. 2 Cs(s) + Br 2 (l) 2 CsBr(s) g. KOH(aq) + HCl(aq) H 2 O(l) + KCl(aq) h. 2 C 2 H 2 (g) + 5 O 2 (g) 4 CO 2 (g) + 2 H 2 O(l) i. Zn(s) + HCl(aq) ZnCl 2 (aq) + H 2 (g) 4

5 Chapter 11 Modern Atomic Theory Topics and question types Waves (wavelength, frequency, electromagnetic spectrum) The Bohr model The wave mechanical model (Schrodinger, orbitals, etc.) Electron configurations Periodic trends (atomic radius and ionization energy) Practice 1. Which of the following is not an example of electromagnetic radiation? a. X-rays b. Visible light c. Sound waves d. UV rays e. All of the above are examples of electromagnetic radiation 2. Which of the following statements about atomic structure is incorrect? a. electrons are located around the nucleus b. when an electron moves to a higher energy level, it is farther from the nucleus c. electrons move in circular orbits around the nucleus d. none of the above are true e. all of the above are true 3. The distance between two crests (or two troughs) in a wave is called the a. frequency b. wavelength c. speed d. amplitude e. color 4. An orbital can be described as a. a path followed by the electron b. the direction in which the electron is moving c. the wavelength of the electron d. an area in space where the electron is likely to be found e. none of the above 5. A photon with a longer wavelength will have a. lower energy b. higher energy c. lower frequency d. higher frequency e. a and c f. b and d 5

6 6. The fact that the emission spectrum of hydrogen consists of lines (rather than a continuous spectrum) led scientists to what conclusion? a. the energy of the hydrogen atom is quantized b. every color of light can be emitted from an excited H atom c. the Bohr model is incorrect d. all of the above 7. How many valence electrons are there in an atom of calcium? a. 1 b. 2 c. 3 d. 4 e How many valence electrons are there in an atom of sulfur? a. 2 b. 4 c. 5 d. 6 e How many unpaired electrons are there in an atom of Se? a. 1 b. 2 c. 3 d. 4 e How many unpaired electrons are there in an atom of cobalt? a. 1 b.2 c. 3 d. 4 e Which of the following orbitals is highest in energy? a. 1s b. 2s c. 3s d. 4s e. 5s 12. Which of the following orbitals is highest in energy? a. 5s b. 5p c. 5d d. 5f e. 5g 13. How many sublevels are there in the fourth energy level (n=4)? a. 1 b. 2 c. 3 d. 4 e Which is NOT a valid orbital designation? a. 6s b. 3p c. 4f d. 2d e. 1s 15. Which of the figures below represents a p orbital? a. b. c. d. e. 16. Which of the figures below represents an s orbital? a. b. c. d. e. 6

7 17. Write the complete electron configuration for magnesium. 18. Write the complete electron configuration for cadmium. 19. Write the complete electron configuration for arsenic (As). 20. Write the complete orbital diagram for carbon. 21. How many unpaired electrons does a carbon atom have? 22. Write the complete orbital diagram for sulfur. 23. How many unpaired electrons does a sulfur atom have? 24. Write the complete orbital diagram for nickel. 25. How many unpaired electrons does an oxygen atom have? 26. In terms of atomic size, which of the following is NOT true? a. F < B < Li b. I > Br > F c. Na < Cs < Fr d. Rb > Sr > Xe e. C > Si > Ge 27. In terms of atomic size, which of the following is true? a. Ca > Cr > K b. Be < Ba < Ra c. B < N < O d. Ne > Ar > Kr e. Fe < Ga < Br 28. In terms of ionization energy, which of the following is true? a. Na > Mg > Al b. Na < K < Rb c. Ba < Ca < Be d. Br < Se < As e. F < O < N 29. Moving down the periodic table, does atomic size increase or decrease? Explain why. 30. Moving across the periodic table (from left to right), does atomic size increase or decrease? Explain why. 31. Moving down the periodic table, does ionization energy increase or decrease? Explain why. 32. Moving across the periodic table (from left to right), does ionization energy increase or decrease? Explain why. 33. How many electrons can there be in one s orbital? a. 1 b. 2 c. 3 d. 4 e How many electrons can there be in the 4d sublevel? a. 2 b. 5 c. 6 d. 10 e Using the shorthand notation (with a noble gas core), write the electron configuration for barium. 36. How many valence electrons does barium have? 37. How many unpaired electrons does barium have? 38. Using the shorthand notation (with a noble gas core), write the electron configuration for Sn. 39. How many unpaired electrons does tin have? 40. What was Erwin Schrodinger s contribution to our understanding of the structure of the atom? 7

8 Chapter 12 Chemical Bonding Topics and question types Electronegativity and polarity Electron configurations of ions Lewis Structures VSEPR (shapes, bond angles) (Honors only) Resonance, bond order Practice 1. Define electronegativity. 2. Put the following elements in order of increasing electronegativity. a) Ca, Ga, Br, K b) C, Sn, Pb, Si 3. Put the following bonds in order of increasing polarity. a) O-S, O-Te, O-Se b) B-F, B-C, B-N, B-B 4. Put the following atoms/ions in order of increasing size. a) P 3-, Cl -, Ar, S 2- b) Mg 2+, Ne, F -, Na + c) Al, Al 3+ d) N, N 3-5. Write an electron configuration for each ion. a) O 2- b) Ca 2+ c) S 2-6. Write an electron configuration for the ion commonly formed by each element. a) K b) Br 7. How many lone pairs of electrons are there in the Lewis structure of SO 2? 8. Draw a Lewis structure for PCl Name the geometry of PCl What is the Cl-P-Cl bond angle in PCl 3? 11. Is PCl 3 polar? 12. Draw a Lewis structure for XeF Draw a Lewis structure for SF Draw a Lewis structure for AlCl Name the geometry for AlCl Is AlCl 3 polar? 17. Draw a Lewis structure for chlorite (ClO 2 - ). Accelerated 18. Draw a Lewis structure for H 2 SO Draw at least 6 resonance structures for SO What is the oxygen-oxygen bond order in ozone (O 3 )? 21. Which has the greater bond length: the nitrogen-oxygen bond in nitrite or the nitrogenoxygen bond in nitrate? Which has the strongest nitrogen-oxygen bonds? Explain. Your explanation must include Lewis structures. 22. Draw several resonance structures for the perchlorate ion. 8

9 ANSWERS Semester 1 Review Guide Chemistry Chapter 2 1. a) 3 b) 4 c) 1 d) 3 e) 3 f) 3 g) 4 2. a) 3.6 x 10 5 b) 4.8 x 10-7 c) x 10 6 d) 7.20 x a) 62,000,000 b) c) a) 18 b) 40. c) 150,000 d) 4, a) 1.44 m b) 4.20 x μg c) 301 K 6. a) ml b) 51 cm g/ml Chapter 3 1. a) P b) C c) P d) C e) P f) C 2. Matter anything that has mass and occupies space (i.e., has volume) Compound a substance composed of 2 or more elements that are chemically bonded together Mixture a combination of 2 or more substances (the components are mixed but are not chemically bonded together) Homogeneous describes a mixture that is uniform throughout. A homogeneous mixture is also called a solution. Heterogeneous describes a mixture that is not uniform throughout. Filtration a process used to separate a mixture consisting of a liquid and an insoluble solid. The mixture is poured over a piece of filter paper, and the solid remains on the filter paper while the liquid (now called the filtrate) passes through. Distillation a process used to separate a mixture of 2 liquids or a liquid and a soluble solid. The substance with the lower boiling point evaporate and then condenses in another flask. See textbook or notes for diagram. Chapter 4 1. Name Relative Mass Charge Location in the atom Proton 1 +1 in the nucleus Neutron 1 0 in the nucleus Electron 1 / outside nucleus (electron cloud) 2. isotopes 3. a) 11 protons, 11 neutrons, 11 electrons; 22 11Na b) 17 protons, 18 neutrons, 18 electrons; 35 17Cl - c) 8 protons, 10 neutrons, 8 electrons; 18 8O d) 15 protons, 16 neutrons, 15 elecctrons; 31 15P 4. a) halogen b) alkali metal c) transition metal d) noble gas e)alkaline earth metal 5. a) Na b) C 6. a) conduct heat and eletrcitiy; b) malleable; c) ductile; d) shiny (lustrous) Chapter 5 1. a) sodium oxide b) nickel(ii) nitrate c) lead(i) oxide d) sulfur hexafluoride e) manganese(ii) hydroxide f) chromium(iii) oxide g) dinitrogen tetrabromide h) nitric acid i) phosphorus pentachloride j) hydrobromic acid 2. a) H 2 SO 4 b) Cu 2 O c) CdCl 2 d) Cs 3 PO 4 e)xef 6 f) K 2 Cr 2 O 7 g) Li 2 S h) Fe 2 (CO 3 ) 3 i) P 2 O 5 j) CrF 2 9

10 Chapter 6 1. a) Cl 2 (g) + 2 KI(aq) I 2 (s) + 2 KCl(aq) b) CaF 2 (s) + H 2 SO 4 (aq) CaSO 4 (s) + 2 HF(g) c) C 5 H 12 (g) + 8 O 2 (g) 5 CO 2 (g) + 6 H 2 O(g) d) PbCl 2 (aq) + K 2 SO 4 (aq) PbSO 4 (s) + 2 KCl(aq) e) Fe(NO 3 ) 3 (aq) + 3 KOH(aq) Fe(OH) 3 (s) + 3 KNO 3 (aq) Chapter 7 1. a) Ba 2+ (aq) + SO 4 2- (aq) BaSO 4 (s) b) H + (aq) + OH - (aq) H 2 O(l) c) Zn(s) + 2 H + (aq) Zn 2+ (aq) + H 2 (g) 2. a) combustion, redox b) combustion, redox, synthesis c) redox, synthesis d) redox, decomposition e) precipitation, double displacement f) redox, synthesis g) acid-base, double displacement h) combustion, redox i) single replacement, redox Chapter C 6. A 11. E 16. B 2. C 7. B 12. E 17. 1s 2 2s 2 2p 6 3s 2 3. B 8. D 13. D 18. 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 6 5s 2 4d D 9. B 14. D 19. 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p 1 x 4p 1 1 y 4p z 5. E 10. C 15. C 20. _ 1s 2s 2p _ _ 1s 2s 2p 3s 3p _ _ 1s 2s 2p 3s 3p 4s 3d E 27. B 10

11 28. C 29. Atomic size increases. Moving down the Periodic Table, additional electrons are added to a new energy level. Since the higher energy levels are farther from the nucleus, the atomic radius increases. 30. Atomic size decreases. Although additional electrons are being added, they are being added to the same energy level and are therefore no farther from the nucleus. Since more protons are being added as well, the positive charge in the nucleus is increasing and therefore pulls the electrons closer to the nucleus. 31. Ionization energy decreases. The highest energy electron is farther from the nucleus and is therefore easier (i.e., requires less energy) to remove. 32. Ionization energy increases. The highest energy electron is closer to the nucleus and is therefore harder (i.e., requires more energy) to remove. 33. B 34. D 35. [Xe] 6s [Kr] 5s 2 4d 10 5p x 1 5p y He treated the electron as a wave and developed the idea of orbitals. Chapter Electronegativity is the tendency of an atom to pull shared electrons towards itself. 2. a) K, Ca, Ga, Br b) Pb, Sn, Si, C 3. a) O-S, O-Se, O-Te b) B-B, B-C, B-N, B-F 4. a) Ar, Cl -, S 2-, P 3- b) Mg 2+, Na +, Ne, F - c) Al 3+, Al d) N, N 3-5. a) 1s 2 2s 2 2p 6 b) 1s 2 2s 2 2p 6 3s 2 3p 6 c) 1s 2 2s 2 2p 6 3s 2 3p 6 6. a) 1s 2 2s 2 2p 6 3s 2 3p 6 b) 1s 2 2s 2 2p 6 3s 2 3p 6 4s 2 3d 10 4p lone pairs trigonal pyramidal 10. approximately 107º (109.5º is acceptable) 11. yes 11

12 trigonal planar 16. no OR Accelerated 19. OR 20. tetrahedral 21. 1½ or 3/2 12

13 22. Nitrate has 3 resonance structures and a N-O bond order of 4/3 or 1 1/3. Nitrite has 2 resonance structures and a N-O bond order of 3/2 or 1 ½. Since a higher bond order indicates a stronger and shorter bond, nitrite (BO 3/2) has stronger N- O bonds, and nitrate (BO 4/3) has longer (weaker) N-O bonds. It is important to understand that both bonds in nitrite are identical and all three bonds in nitrate are identical