Practice Problems- Set# 5 Text book Problems
Practice Problems- Set# 5 Text book Problems: 8.1, 8.2, 8.3, 8.4, 8.5, 8.13,8.14, 8.15, 8.16,, 8.17, 8.19, 8.21, 8.22,, 8.23,, 8.24, 8.27, 8.28, , 8.29, 8.31, 8.32, 8.33, 8.35, 8.36, 8.37, 8.39, 8. 53, 8.54, 8.55, 8.56, 8.59, , 8.60, 8.61, 8.62, 8.67, 8.73, 8.74, 8.75, 8.89, 8.90, 8.92 Review extra practice exercises ( 8.1….-8.10; pg 337 for answers) I)
Short Answer Questions/Fill in the blanks:
1) Using shorthand notation, the the ground-state electron configuration for Sr 2+ is predicted to be_____
2) Using shorthand notation, the the ground-state electron configuration for C 4 is predicted to be_____.
3) The ionic radius of Cs + is _____ than the atomic radius of Cs, and the ionic radius of I is _____ than the atomic radius of I.
4) The element in period 4 with the smallest first ionization energy is _____.
5) The element in period 3 with the smallest seventh ionization energy is _____.
6) The element in group 7A with the least favorable (least negative) electron affinity is _____.
7) Lattice energy increases with _____ cation and anion charges and _____ cation and anion radii.
8) The group 4A element that always obeys the octet rule in its stable compounds is _____.
II)Study Questions… Rutherford’s atom 1. What is meant by the expression “nuclear atom”? 2. Although Rutherford’s experiments confirmed the existence of the positively charged nucleus, they left many questions about the structure of the atom unanswered. Discuss. Electromagnetic Radiation 3. What is electromagnetic radiation? At what speed does electromagnetic radiation travel? 4. How are different types of electromagnetic radiation similar? How do they diffr? 5. What is a “packet” of electromagnetic radiation called? 6. What is white light? Colored light? 7. What is the relationship between wavelength and the amount of energy carried by its photons? 8. Explain, in termsof absorbed and reflected light,why a blue object appears blue. 9. What produces gamma rays? How are X-rays used? 10. Why do microwave ovens heat food but not the dish the food is on? 11. Why should excess exposure to ultraviolet light, gamma rays and X-rays be avoided? 12. What type of electromagnetic radiation is used in communiation devices such as cellular telephones? Emission of energy by atoms/Bohr model 13. What is observed when salts containing metal ions sucha s Na+, Cu2+ or Li+ are heated ina flame? 14. What is an emission spectrum? Use the Bohr model to explain why the emission spectra of atoms consist of distinct lines at specific wavelengths. 15. Describe briefly why the study of electromagnetic radiation has been important to our understanding of the arrangement of electrons in atoms. 16. What does it mean to say that the hydrogen atom has “discrete” energy levels? How is this fact reflected in the radiation that excited hydrogen atoms emit? 17. What are the essential point sof Bohr’s theory of the structur eof the hydrogen atom? 18. Why was Bohr’s theory for the hydrogen atom accepted initially, and why was it ultimately discarded? 19. Explain the difference between a Bohr orbit and a quantum mechanical orbital. 20. Two of the emission wavelengths in the hydrogen emission spectrum are 410 nm and 434 nm. One of these is due to the n=6 to n=2 transition and the other is due to the n=5 to n=2 transition. Which wavelenght goes with which transition? Wave Mechanical Model/Hydrogen Orbitals 21. What major assumptions ( analogous to that already demonstrated for electromagnetic radiation) did de Broglie and Schrodinger make about the motion of tiny particles? 22. When describing electrons ina na orbital we use arrows pointing upwarda dn downward to indiate what property? What is Pauli exclusion principle? Why is it important when writing electron configurations? 23. What is Hund’s rule? Why is it important when writing electron configurations? 24. Give some examples of the explanatory power of the quantum mechanical model.
25. Which electron is, on an average close to the nucleus: an electron in a 2s orbital or an electron in a 3s orbital? Which electron is, on an average further from the nucleus: an electron in a 3p orbital or a 4p orbital? Electron configurations and the Periodic Table 26. Why do we believe that the valence electrons of calcium and potassium reside in the 4s orbital rather than in the 3d orbital? 27. Would you expect the valence electrons of rubidiuma nd strontium to reside in the 5s or the 4d orbital? Why? 28. Write electron configurations for each of the following transition metals: i)Co ii) Ni iii) Cu iv) Zn 29. Write the complete orbital diagram for each of the following elements, using boxes to represent orbitals and arrows to represent electrons. i) Aluminum, Z = 13; ii) phosphorous, Z = 15; iii) bromine, Z =35, iv)Argon, Z =18 30) How many valence electrons does each of the following atoms possess? i) Cs, B, Ba, Se 31. Write electron configurations for the following ions: Ca2+, S2-, Al3+ Atomic Properties and the Periodic Table 32. In each of the following sets of elements, which element would be expected to have the highest ionization energy? a) Cs, K, Li b) Ba, Sr, Ca c) I, Br, Cl d) Mg, Si, S 33. In each of the following sets of elements, indicate which element has the smallest atomic size. a) Na, K, Rb b) Na, Si, S c) N, P, As d) N, O, F 34. One bit of evidence that the present theory of atomic structure is “correct” lies in the magnetic properties of matter. Atoms with unpaired electrons are attracted by magnetic fields and thus aresaid to exhibit paramagnetism.The degree to which this effect is observed is directly related to the number of unpaired electrons present in the atom. On the basis of the electron orbital diagrams for the following elements, indicate which atoms would be expected to be paramagnetic, and tell how many unpaired electrons each atom contains: a) phosphorous, Z =15 b) iodine, Z = 53 c) germanium, Z = 32
III) Answer the following in a few lines.. Interpreting the energies of characterisitic atomic transitions 1. Sodium street lamps give off a characteristic yellow light of wavelength 588nm. What is the frequency of this light? What is the energy per mole (in kJmol-1) of these photons? 2. What is the energy per mole of the 514 nm photons emitted in the Ar+ transition which is
exploited to make green laser light?
3.
4. 5. 6.
Accounting for the variation in electron affinity Account for the large decrease in electron affinity between lithium and beryllium despite the increase in nuclear charge. Accounting for the variation in ionization energy Account for the decrease in first ionization energy between phosphorous and sulfur. Account for the decrease in first ionization energy between fluorine and chlorine. Accounting for trends in effective nuclear charge Suggest a reason why the increase in Zeff for a 2p electron is smaller between N and O than between C & N given that the configurations of the three atoms are: C: [He] 2s22p2 N: [He]2s22p3 O: [He]2s22p4
7. Account for the larger increase in effective nuclear charge for a 2p electron on going from B to C compared to a 2s electron on going from Li to Be. Probablity of being close to the nucleus 8. Which orbital, 3p or 3d, gives an electron a greater probability of being close to the nucleus? Additional…. 9. How do we know that the energy levels of the hydrogen atom are not continuous, as physicsts originally assumed? 10. Excessive exposure to sunlight increases the risk skin cancer because some of the photons have enough energy to breakchemical bonds in biological molecules. These bond require approximately 250-800 kJ/mole of energy to break. The energy of a single photon is given by E =hc/ where E is the energy of the photon in J, h is Planck’s constant ( 6.626 x 10 -34 Js) and c is the speed of light ( 3.00 x 108 m/s). Determine which of the following kinds of light contain enough energy to break chemical bonds in biological molecules by calculating the total energy in 1mol of photons for light of each wavelength. a) Infra red light ( 1500 nm) b) Visible light ( 500 nm) c) Ultraviolet light ( 150 nm)
Short Answer Questions/Fill in the blanks:
1) Using shorthand notation, the the ground-state electron configuration for Sr 2+ is predicted to be_____
2) Using shorthand notation, the the ground-state electron configuration for C 4 is predicted to be_____.
3) The ionic radius of Cs + is _____ than the atomic radius of Cs, and the ionic radius of I is _____ than the atomic radius of I.
4) The element in period 4 with the smallest first ionization energy is _____.
5) The element in period 3 with the smallest seventh ionization energy is _____.
6) The element in group 7A with the least favorable (least negative) electron affinity is _____.
7) Lattice energy increases with _____ cation and anion charges and _____ cation and anion radii.
8) The group 4A element that always obeys the octet rule in its stable compounds is _____.
II)Study Questions… Rutherford’s atom 1. What is meant by the expression “nuclear atom”? 2. Although Rutherford’s experiments confirmed the existence of the positively charged nucleus, they left many questions about the structure of the atom unanswered. Discuss. Electromagnetic Radiation 3. What is electromagnetic radiation? At what speed does electromagnetic radiation travel? 4. How are different types of electromagnetic radiation similar? How do they diffr? 5. What is a “packet” of electromagnetic radiation called? 6. What is white light? Colored light? 7. What is the relationship between wavelength and the amount of energy carried by its photons? 8. Explain, in termsof absorbed and reflected light,why a blue object appears blue. 9. What produces gamma rays? How are X-rays used? 10. Why do microwave ovens heat food but not the dish the food is on? 11. Why should excess exposure to ultraviolet light, gamma rays and X-rays be avoided? 12. What type of electromagnetic radiation is used in communiation devices such as cellular telephones? Emission of energy by atoms/Bohr model 13. What is observed when salts containing metal ions sucha s Na+, Cu2+ or Li+ are heated ina flame? 14. What is an emission spectrum? Use the Bohr model to explain why the emission spectra of atoms consist of distinct lines at specific wavelengths. 15. Describe briefly why the study of electromagnetic radiation has been important to our understanding of the arrangement of electrons in atoms. 16. What does it mean to say that the hydrogen atom has “discrete” energy levels? How is this fact reflected in the radiation that excited hydrogen atoms emit? 17. What are the essential point sof Bohr’s theory of the structur eof the hydrogen atom? 18. Why was Bohr’s theory for the hydrogen atom accepted initially, and why was it ultimately discarded? 19. Explain the difference between a Bohr orbit and a quantum mechanical orbital. 20. Two of the emission wavelengths in the hydrogen emission spectrum are 410 nm and 434 nm. One of these is due to the n=6 to n=2 transition and the other is due to the n=5 to n=2 transition. Which wavelenght goes with which transition? Wave Mechanical Model/Hydrogen Orbitals 21. What major assumptions ( analogous to that already demonstrated for electromagnetic radiation) did de Broglie and Schrodinger make about the motion of tiny particles? 22. When describing electrons ina na orbital we use arrows pointing upwarda dn downward to indiate what property? What is Pauli exclusion principle? Why is it important when writing electron configurations? 23. What is Hund’s rule? Why is it important when writing electron configurations? 24. Give some examples of the explanatory power of the quantum mechanical model.
25. Which electron is, on an average close to the nucleus: an electron in a 2s orbital or an electron in a 3s orbital? Which electron is, on an average further from the nucleus: an electron in a 3p orbital or a 4p orbital? Electron configurations and the Periodic Table 26. Why do we believe that the valence electrons of calcium and potassium reside in the 4s orbital rather than in the 3d orbital? 27. Would you expect the valence electrons of rubidiuma nd strontium to reside in the 5s or the 4d orbital? Why? 28. Write electron configurations for each of the following transition metals: i)Co ii) Ni iii) Cu iv) Zn 29. Write the complete orbital diagram for each of the following elements, using boxes to represent orbitals and arrows to represent electrons. i) Aluminum, Z = 13; ii) phosphorous, Z = 15; iii) bromine, Z =35, iv)Argon, Z =18 30) How many valence electrons does each of the following atoms possess? i) Cs, B, Ba, Se 31. Write electron configurations for the following ions: Ca2+, S2-, Al3+ Atomic Properties and the Periodic Table 32. In each of the following sets of elements, which element would be expected to have the highest ionization energy? a) Cs, K, Li b) Ba, Sr, Ca c) I, Br, Cl d) Mg, Si, S 33. In each of the following sets of elements, indicate which element has the smallest atomic size. a) Na, K, Rb b) Na, Si, S c) N, P, As d) N, O, F 34. One bit of evidence that the present theory of atomic structure is “correct” lies in the magnetic properties of matter. Atoms with unpaired electrons are attracted by magnetic fields and thus aresaid to exhibit paramagnetism.The degree to which this effect is observed is directly related to the number of unpaired electrons present in the atom. On the basis of the electron orbital diagrams for the following elements, indicate which atoms would be expected to be paramagnetic, and tell how many unpaired electrons each atom contains: a) phosphorous, Z =15 b) iodine, Z = 53 c) germanium, Z = 32
III) Answer the following in a few lines.. Interpreting the energies of characterisitic atomic transitions 1. Sodium street lamps give off a characteristic yellow light of wavelength 588nm. What is the frequency of this light? What is the energy per mole (in kJmol-1) of these photons? 2. What is the energy per mole of the 514 nm photons emitted in the Ar+ transition which is
exploited to make green laser light?
3.
4. 5. 6.
Accounting for the variation in electron affinity Account for the large decrease in electron affinity between lithium and beryllium despite the increase in nuclear charge. Accounting for the variation in ionization energy Account for the decrease in first ionization energy between phosphorous and sulfur. Account for the decrease in first ionization energy between fluorine and chlorine. Accounting for trends in effective nuclear charge Suggest a reason why the increase in Zeff for a 2p electron is smaller between N and O than between C & N given that the configurations of the three atoms are: C: [He] 2s22p2 N: [He]2s22p3 O: [He]2s22p4
7. Account for the larger increase in effective nuclear charge for a 2p electron on going from B to C compared to a 2s electron on going from Li to Be. Probablity of being close to the nucleus 8. Which orbital, 3p or 3d, gives an electron a greater probability of being close to the nucleus? Additional…. 9. How do we know that the energy levels of the hydrogen atom are not continuous, as physicsts originally assumed? 10. Excessive exposure to sunlight increases the risk skin cancer because some of the photons have enough energy to breakchemical bonds in biological molecules. These bond require approximately 250-800 kJ/mole of energy to break. The energy of a single photon is given by E =hc/ where E is the energy of the photon in J, h is Planck’s constant ( 6.626 x 10 -34 Js) and c is the speed of light ( 3.00 x 108 m/s). Determine which of the following kinds of light contain enough energy to break chemical bonds in biological molecules by calculating the total energy in 1mol of photons for light of each wavelength. a) Infra red light ( 1500 nm) b) Visible light ( 500 nm) c) Ultraviolet light ( 150 nm)
