7. How many unpaired electrons are there in an atom of tin in its ground state? 2

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1 Name period AP chemistry Unit 2 worksheet 1. List in order of increasing energy: 4f, 6s, 3d,1s,2p 1s, 2p, 6s, 4f 2. Explain why the effective nuclear charge experienced by a 2s electron in boron is greater than that for the 2p electron. The 2p electron in boron is shielded from the full charge of the nucleus by the 2s electrons, so the 2p electron experiences a smaller effective nuclear charge 3. Explain why the effective nuclear charge experienced by a 2s electron in aluminum is greater than that for the 2s electron experienced by boron. Aluminum has more protons than boron so the 2s electron will experience a greater effective nuclear charge 4. Which should experience the greater nuclear charge, a 2p electron in oxygen or a 2p electron in neon? A 2p electron in neon experiences a greater effective nuclear charge 5. How many f orbitals have n=3? 0 6. Two electrons in an atom both occupy the 1s orbital. What quantity must be different for the two electrons? They must have opposite spins 7. How many unpaired electrons are there in an atom of tin in its ground state? 2 8. Of the following elements, which one is most likely to form an ion through the loss of two electrons? a. strontium b. sulfur c. sodium d. chlorine e. aluminum 9. An atom has two electrons with principal quantum number (n) = 1, eight electrons with principal quantum number (n) = 2 and seven electrons with principal quantum number (n) = 3. From this data, supply the following values (if insufficient information is given, say so). (a) The mass number. not enough info (b) The atomic number. 17 (c) The electron configuration. _1s 2 2s 2 2p 6 3s 2 3p What is the maximum number of electrons that can occupy each of the following a. 3d 10 b. 4s 2 c. 2 nd shell 8 d. n=3 18 a. 2p 6 f. 5f 14 g. One 2p orbital 2 h. n= Write the orbital notation (can use noble gas) for each of the following a. Sc b. Si [Ar] 4s 3d [Ne] 3s 3p c. Sn d. Mn _ _ _ _ _ [Kr] 5s 4d 5p [Ar] 4s 3d

12. Write the noble gas configuration for the following a. Rb b. Se c. Zn [Kr]5s 1 [Ar]4s 2 3d 10 4p 4 [Ar]4s 2 3d 10 d. Pb e. Mn f. N [Xe]6s 2 5d 10 4f 14 6p 2 [Ar]4s 2 3d 5 [He]2s 2 2p 3 13.

(3) (a) [Ne] 3s 2 3p 2 Si (b) [Ar] 4s 2 3d 7 Co (c) [Xe] 6s 2 Ba 15. Which element could be represented by the complete PES spectrum to the right? A. Li B. B C. N D. Ne 16.

2 12. Write the noble gas configuration for the following a. Rb b. Se c. Zn [Kr]5s 1 [Ar]4s 2 3d 10 4p 4 [Ar]4s 2 3d 10 d. Pb e. Mn f. N [Xe]6s 2 5d 10 4f 14 6p 2 [Ar]4s 2 3d 5 [He]2s 2 2p Write the full electronic configuration for argon 1s 2 2s 2 2p 6 3s 2 3p Identify the element from the electron configurations of atoms shown below. (3) (a) [Ne] 3s 2 3p 2 Si (b) [Ar] 4s 2 3d 7 Co (c) [Xe] 6s 2 Ba 15. Which element could be represented by the complete PES spectrum to the right? A. Li B. B C. N D. Ne 16. Which of the following best explains the relative positioning and intensity of the 2s peaks in the following spectra? a)be has a greater nuclear charge than Li and more electrons in the 2s orbital b)be electrons experience greater electron-electron repulsions than Li electrons c)li has a greater pull from the nucleus on the 2s electrons, so they are harder to remove d)li has greater electron shielding by the 1s orbital, so the 2s electrons are easier to remove 17. Which will be closer to the nucleus, the n=3 electron shell in Ar or the n=3 shell in Kr? Kr 18. Arrange the following atoms in order of increasing atomic radius: F, P, S, As and explain why. F, S, P, As (atoms get bigger when a shell is added or less protons in the same number of shells) 19. Arrange the following atoms in order of increasing atomic radius: Al, Nb, Se, F, Mn and explain why. F, Al, Se,Mn, Nb, (atoms get bigger when a shell is added or less protons in the same number of shells) 20. An element having the configuration [Xe]6s 1 belongs to the group: a. alkaline earth metals b. alkali metals c. halogens d. noble gases e. none of these 21. Explain in terms of electron configurations, why hydrogen exhibits properties similar to both lithium and fluorine It has one valence electron like lithium, but only needs one valence electron to have a full shell like fluorine 22. Which of the following statements are true a. All are false b. the first ionization energy of fluorine is greater than the first ionization energy of oxygen c. as the atomic number increases within a group of the main group elements, the tendency is for first ionization energy to increase d. it is easier to remove an electron from Na + than from Na. e. all particles with the electron configuration [Ar]4s 2 have the same ionization energy.

3 23. Consider the element Scandium, atomic # 21. (a) If the electronic configuration of the element were constructed "from scratch", into which orbital (and into which shell) would the final electron be placed? 3d (b) When scandium forms an ion with a charge of +1, from which orbital (and from which shell) would the electron be removed? 4s 24. Based on their position on the periodic table, predict which atom of the following pairs will have the largest first ionization energies. In each case explain with electron configuration and effective nuclear charge a. O, Ne b. Mg, Sr c. K, Cr d. Br, Sb e. Ga, Ge Ne Mg Cr Br Ge Smaller atoms have a harder time losing their electrons because the valence electrons feel a greater effective nuclear charge so it takes more energy to remove the electron. 25. Identify two positive and two negative ions that are isoelectronic with argon. (4) (a) Two Positive ions _K + Ca 2+ (b) Two Negative ions Cl - S Compare the elements sodium and magnesium with respect to the following properties a. Electron configuration 1s 2 2s 2 2p 6 3s 1 ; 1s 2 2s 2 2p 6 3s 2 b. Most common ionic charge+1; +2 c. First ionization energy magnesium has a higher first ionization energy d. Atomic radius magnesium is smaller than sodium 27. Compare the elements fluorine and chlorine with respect to the following properties a. Electron configuration b. Most common ionic charge c. Atomic radius 1s 2 2s 2 2p 5 ; 1s 2 2s 2 2p 6 3s 2 3p 5-1;-1 Chlorine is larger than fluorine 28. Write the noble gas configuration a. Fe 3+ b. Ni 2+ [Ar]3d 5 [Ar]3d Arrange the atoms and ions in each of the following sets in order of increasing size a. Br -, Na +, Mg 2+ b. Ar, Cl -, S 2-, K + Mg, Na, Br K, Ar, Cl, S 30. Using the periodic table, select the most electronegative atom in each of the following sets a. B, Be, C, Si b. Zn, Ga, Ge, As c. Na, Mg, K, Ca C As Mg 31. How many protons, neutrons, and electrons are in the following a. 65 Zn 2+ b. 40 Ar c. 14 N 3- d. 23 Na + 30,35,28 18,22,18 7,7,10 11,12, Which ions are cations in the previous problem, which are anions? Cations: A,d anions: c 33. How many valence electrons does each of the following atoms have? a. C b. Ca c. H d. Pb e. Ar f. Cl

4 34. The ionization energies for an element are listed below First second third fourth fifth 8 ev 15eV 80eV 109eV 141 ev Based on the ionization energies, the element is most likely to be a. Sodium b. magnesium c. aluminum d. silicon e. phosphorus 35. Which of the following contains only atoms that are diamagnetic in their ground state? a. Kr, Ca, and P b. Ne, Be, and Zn c. Ar, K, and Ba d. He, Sr, and C 36. Which of the following is the electron configuration of an excited atom that is likely to emit a quantum of energy? (A) 1s 2 2s 2 2p 6 3s 2 3p 1 (B) 1s 2 2s 2 2p 6 3s 2 3p 5 (C) 1s 2 2s 2 2p 6 3s 2 (D) 1s 2 2s 2 2p 6 3s 1 3p The bond energy of fluorine in 159 kj mol -1. i. Determine the energy, in J, of a photon of light needed to break one F-F bond x J ii. Determine the frequency of this photon in s x s -1 iii. Determine the wavelength of this photon in nanometers 750 nm 38. Barium imparts a characteristic green color to a flame. The wavelength of this light is 551 nm. Determine the energy involved in kj/mol 220 kj/mol 39. Explain the difference between metallic, ionic, and covalent bonding Metallic- cations share a sea of electrons Ionic- atoms give and take electrons Covalent- atoms share electrons 40. Why can metals conduct electricity? electrons are free to move around 41. Label each of the following as ionic, metallic, or covalent a. NaOH b. N 2O c. KCl d. HF e. O 2 f. Al foil ionic covalent ionic covalent covalent metallic 42. Which of the following forms molecules? a. K 2CO 3 b. F 2 c. BaCl 2 d. H 2O e. Fe 2O Predict the chemical formula of the ionic compound formed between the following pairs of elements a. Al and F b. K and S c. Mg and N d. Ba and O AlF 3 K 2S Mg 3N 2 BaO 44. Arrange the following substances according to their expected lattice energies, listed them from lowest lattice energy to highest: LiCl, KCl, KBr, CaO KBr< KCl < LiCl < CaO 45. Explain the following trends in lattice energy a. MgO > MgS b. LiF > CsBr c. CaO > KF Oxygen is smaller than S Cs is larger than Li CaO has larger charges than KF Br is larger than F 46. How is bonding in Cl 2 different than NaCl? In Cl 2, the electrons are being shared between the two chlorine atoms, but in NaCl, the Cl is taking the electron from the sodium

5 47. Draw the Lewis structure for O 2. The bond in O 2 is shorter than the O-O single bond. Explain this observation. O=O A double bond is shorter than a single bond 48. Predict whether the following compounds are molecular or ionic a. B 2H 6 b. CH 3OH c. LiNO 3 d. Sc 2O 3 e. CsBr f. NOCl g. Ag 2SO 4 Molecular molecular ionic ionic ionic covalelnt ionic 49. Which of the following bonds are polar? a. P-O b. S-F c. Br-Br 50. Arrange the bonds in order of increasing polarity a. C-F, O-F, Be-F b. N-Br, P-Br, O-Br O-F, C-F, Be-F P-Br, N-Br, O-Br 51. Label each compound as ionic, polar covalent, or nonpolar covalent a. CO b. MgO c. Cl 2 d. AlF 3 Polar covalent ionic nonpolar covalent ionic 52. Draw the Lewis structure for the following see in class a. CO b. N 2 c. SF 2 - d. ClO 2 e. PCl 3 e. H 2CO (both H bonded to C) 53. A.Draw the Lewis electron-dot structures for CO 3 2-, CO 2 and CO, including resonance structures where appropriate. See in class b. Which of the three species has the shortest C-O bond length? Explain the reason for your answer. CO carbon monoxide forms triple bonds which are shorter than double in CO 2 and the 1 1/3 bond in CO 3 c. Account for the fact that the carbon-oxygen bond length in CO 3 2 is greater than the carbon-oxygen bond length in CO /3 bond is longer than a double bond 54. How can the concept of resonance be used to explain that all six C-C bonds in benzene are equal in length. The 3 double bonds are being shared between all 6 atoms so the bond order for each C-C bond is Use simple structure and bonding models to account for each of the following: (a) The bond length between the two carbon atoms is shorter in C 2H 4 than in C 2H 6. C 2H 4 forms double bonds and C2H6 forms single bonds. Double bonds are shorter than single (b) All the bond lengths in SO 3 are identical and are shorter than a sulfur-oxygen single bond. SO 3 forms resonance so the bonds are each 1 1/3 which is shorter than a single 56. Draw the Lewis structure for each of the following molecules see in class a. CO 3 2- b. BH 3 c. I 3 - d. XeF 4 e. AsF In the Lewis structure for CH2Cl2, what is the number of unshared electron pairs a. 2 b. 8 c. 10 d. 6 e Which one of the following molecules contains a triple bond? a. PF 3 b. NF 3 c. C 2H 2 d. H 2CO e. HOF

6 59. Which of the following has the greatest dipole moment? a. H 2 b. HCl c. HF d. CO 60. Which of the following will conduct electricity? (there can be more than one answer) a. solid Mg b. solid NaCl c. aqueous MgCl 2 d. liquid Sn e. solid CO 2 f. liquid N Predict which of the following will have the highest boiling point a. CO 2 b. N 2 c. NaCl 62. Which of the following ionic compounds has the smallest lattice energy? a. Na 2O b. LiF c. CaO d. CaCl 2 e. MgS 63. Bronze (Cu and Sn); Steel (Fe and C). Which of the following correctly describes the malleability of both alloys compared to their primary metal? a. Bronze s malleability would be comparable to that of copper, but steel s malleability would be significantly lower than that of iron. b. Bronze s malleability would be significantly higher than that of copper, but steel s malleability would be comparable to that of iron. c. Both bronze and steel would have malleability values similar to those of their primary metals d. Both bronze and steel would have malleability values greater than those of their primary metal. Review 64. Which of the following species contain more electrons than neutrons? a. 2 H b. 11 B c. 16 O 2- d. 19 F How many protons, neutrons, and electrons are in an 56 26Fe atom? Protons Neutrons Electrons a b c d Calculate the following to the correct number of significant figures. a.1.27g/5.296cm 3 b. 12.2mL mL c g g d. 0.1 m x 3.21m g/ml 12.6 ml g 0.3 m The density of pure silver is 10.5 g/cm 3. If 5.25 g of pure silver pellets are added to a graduated cylinder containing 11.2 ml of water, to what level will the water in the cylinder rise? 11.7 ml 68. Let s pretend you are holding two atoms of carbon that are isotopes. Describe what the two atoms have in common and what they have different. They have the same number of protons, but one would have more mass than the other one. 69. what is the mass, in grams, of 1.75 x molecules of caffeine, C 8H 10N 4O 2? g 70. Determine the empirical formula of the compound with the following compositions by mass 10.4 percent C, 27.8 percent S, and 61.7 percent Cl CSCl A 2.00g sample of limestone was dissolved in hydrochloric acid and all the calcium present in the sample was converted to Ca 2+ (aq). Excess ammonium oxalate solution, (NH 4) 2C 2O 4(aq), was added to the solution to precipitate the calcium ions as calcium oxalate, CaC 2O 4(s). The precipitate was filtered, dried and weighed to a constant mass of 2.43g. Determine the percentage by mass of calcium in the limestone sample. 38%

7 72. Naphthalene is a hydrocarbon that is used for moth balls. A sample was burned in pure oxygen to produce g of carbon dioxide and grams of water. If the molecular mass of the compound is 128 g/mole so what is the empirical formula and molecular formula for this compound? C 5H 4 and C 10H 8

Test Review # 5. Chemistry: Form TR5-8A. Average Atomic Mass. Subatomic particles.

Test Review # 5. Chemistry: Form TR5-8A. Average Atomic Mass. Subatomic particles. Chemistry: Form TR5-8A REVIEW Name Date Period Test Review # 5 Subatomic particles. Type of Particle Location Mass Relative Mass Charge Proton Center 1.67 10-27 kg 1 +1 Electron Outside 9.11 10-31 kg 0-1