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7.E: Electrons and Chemical Bonds (Exercises)

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    9.3: The Electromagnetic Spectrum

    1. Choose the correct word for the following statement. Blue light has a (longer or shorter) wavelength than red light.
    2. Choose the correct word for the following statement. Yellow light has a (higher or lower) frequency than blue light.
    3. Choose the correct word for the following statement. Green light has a (larger or smaller) energy than red light.
    4. If "light A" has a longer wavelength than "light B", then "light A" has _______________ "light B".
      (a) a lower frequency than
      (b) a higher frequency than
      (c) the same frequency as
    5. If "light C" has a shorter wavelength than "light D", then "light C" has _______________ "light D".
      (a) a larger energy than
      (b) a smaller energy than
      (c) the same energy as
    6. If "light E" has a higher frequency than "light F", then "light E" has __________________ "light F".
      (a) a longer wavelength than
      (b) a shorter wavelength than
      (c) the same wavelength as
    7. If "light G" has a higher frequency than "light H", then "light G" has __________________ "light H".
      (a) a larger energy than
      (b) a smaller energy than
      (c) the same energy as
    8. If "light J" has larger energy than "light K", then "light J" has __________________ "light K".
      (a) a shorter wavelength than
      (b) a longer wavelength than
      (c) the same wavelength as
    9. Which of the following statements is true?
      (a) The frequency of green light is higher than the frequency of blue light and the wavelength of green light is longer than the wavelength of blue light.
      (b) The frequency of green light is higher than the frequency of blue light and the wavelength of green light is shorter than the wavelength of blue light.
      (c) The frequency of green light is lower than the frequency of blue light and the wavelength of green light is shorter than the wavelength of blue light.
      (d) The frequency of green light is lower than the frequency of blue light and the wavelength of green light is longer than the wavelength of blue light.
      (e) The frequency of green light is the same as the frequency of blue light and the wavelength of green light is shorter than the wavelength of blue light.
    10. As the wavelength of electromagnetic radiation increases:
      (a) its energy increases.
      (b) its frequency increases.
      (c) its speed increases.
      (d) more than one of the above statements is true.
      (e) none of the above statements is true.
    11. List three examples of electromagnetic waves.
    12. Why do white objects appear white?
    13. Name the colors present in white light in order of increasing frequency.
    14. Why do objects appear black?

    9.4: The Bohr Model: Atoms with Orbits

    1. Decide whether each of the following statements is true or false:
      (a) Niels Bohr suggested that the electrons in an atom were restricted to specific orbits and thus could only have certain energies.
      (b) Bohr's model of the atom can be used to accurately predict the emission spectrum of hydrogen.
      (c) Bohr's model of the atom can be used to accurately predict the emission spectrum of neon.
      (d) According to the Bohr model, electrons have more or less energy depending on how far around an orbit they have traveled.
    2. According to the Bohr model, electrons in an atom can only have certain, allowable energies. As a result, we say that the energies of these electrons are _______.
    3. The Bohr model accurately predicts the emission spectra of atoms with…
      (a) less than 1 electron.
      (b) less than 2 electrons.
      (c) less than 3 electrons.
      (d) less than 4 electrons.
    4. Consider an He+ atom. Like the hydrogen atom, the He+ atom only contains 1 electron, and thus can be described by the Bohr model. Fill in the blanks in the following statements.
      (a) An electron falling from the n = 2 orbit of He+ to the n = 1 orbit of He+ releases ______ energy than an electron falling from the n = 3 orbit of He+ to the n = 1 orbit of He+.
      (b) An electron falling from the n = 2 orbit of He+ to the n = 1 orbit of He+ produces light with a ______ wavelength than the light produced by an electron falling from the n = 3 orbit of He+ to the n = 1 orbit of He+.
      (c) An electron falling from the n = 2 orbit of He+ to the n = 1 orbit of He+ produces light with a ______ frequency than the light produced by an electron falling from the n = 3 orbit of He+ to the n = 1 orbit of He+.
    5. According to the Bohr model, higher energy orbits are located (closer to/further from) the atomic nucleus. This makes sense since negative electrons are (attracted to/repelled from) the positive protons in the nucleus, meaning it must take energy to move the electrons (away from/towards) the nucleus of the atom.
    6. According to the Bohr model, what is the energy of an electron in the first Bohr orbit of hydrogen?
    7. According to the Bohr model, what is the energy of an electron in the tenth Bohr orbit of hydrogen?
    8. According to the Bohr model, what is the energy of an electron in the seventh Bohr orbit of hydrogen?
    9. If an electron in a hydrogen atom has an energy of −6.06×10−20 J, which Bohr orbit is it in?
    10. If an electron in a hydrogen atom has an energy of −2.69×10−20 J, which Bohr orbit is it in?
    11. If an electron falls from the 5th Bohr orbital of hydrogen to the 3rd Bohr orbital of hydrogen, how much energy is released (you can give the energy as a positive number)?
    12. If an electron falls from the 6th Bohr orbital of hydrogen to the 3rd Bohr orbital of hydrogen, what wavelength of light is emitted? Is this in the visible light range?

      9.5: Electron Configurations and the Periodic Table

    13. Use the Periodic Table to determine the energy level of the valence electrons in each of the following elements.
      (a) B
      (b) Ga
      (c) Rb
      (d) At
      (e) He
    14. Fill in the blanks:
      (a) B is in the __ level block of the Periodic Table
      (b) Sr is in the __ level block of the Periodic Table
      (c) Fe is in the __ level block of the Periodic Table
      (d) Cs is in the __ level block of the Periodic Table
      (e) O is in the __ level block of the Periodic Table
    15. Use the Periodic Table to determine the energy level and sublevel of the highest energy electrons in each of the following elements:
      (a) N
      (b) Ca
      (c) Rb
      (d) P
      (e) In
    16. Decide whether each of the following statements is true or false.
      (a) Li has valence electrons in the n = 1 energy level.
      (b) Si has valence electrons in the n = 3 energy level.
      (c) Ga has valence electrons in the n = 3 energy level.
      (d) Xe has valence electrons in the n = 5 energy level.
      (e) P has valence electrons in the n = 2 energy level.
    17. Match the element to the sublevel block it is found in:
      (a) C i. s sublevel block
      (b) Cs ii. p sublevel block
      (c) Ce iii. d sublevel block
      (d) Cr iv. f sublevel block
    18. The first row of the Periodic Table has:
      (a) 1 element
      (b) 2 elements
      (c) 3 elements
      (d) 4 elements
      (e) 5 elements
    19. Use the Periodic Table to determine which of the following elements has the highest energy valence electrons.
      (a) Sr
      (b) As
      (c) H
      (d) At
      (e) Na
    20. Use the Periodic Table to determine which of the following elements has the lowest energy valence electrons.
      (a) Ga
      (b) B
      (c) Cs
      (d) Bi
      (e) Cl
    21. Which energy level does the first row in the d sublevel block correspond to?

    22. Identify the element with each ground state electron configuration.

      1. [He]2s22p1
      2. [Ar]4s23d8
      3. [Kr]5s24d105p4
      4. [Xe]6s2

    23. Identify the element with each ground state electron configuration.

      1. [He]2s22p1
      2. [Ar]4s23d8
      3. [Kr]5s24d105p4
      4. [Xe]6s2

    24. Give the complete electron configuration for each element.

      1. magnesium
      2. potassium
      3. titanium
      4. selenium
      5. iodine
      6. uranium
      7. germanium

    25. Give the complete electron configuration for each element.

      1. tin
      2. copper
      3. fluorine
      4. hydrogen
      5. thorium
      6. yttrium
      7. bismuth

    26. Write the valence electron configuration for each element:

      1. samarium
      2. praseodymium
      3. boron
      4. cobalt

    9.6: Lewis Electron Dot Diagrams

    1. Explain why the first two dots in a Lewis electron dot diagram are drawn on the same side of the atomic symbol.
    2. Is it necessary for the first dot around an atomic symbol to go on a particular side of the atomic symbol?
    3. What column of the periodic table has Lewis electron dot diagrams with two electrons?
    4. What column of the periodic table has Lewis electron dot diagrams that have six electrons in them?
    5. Draw the Lewis electron dot diagram for each element.
      1. strontium
      2. silicon
    6. Draw the Lewis electron dot diagram for each element.
      1. krypton
      2. sulfur
    7. Draw the Lewis electron dot diagram for each element.
      1. titanium
      2. phosphorus
    8. Draw the Lewis electron dot diagram for each element.
      1. bromine
      2. gallium
    9. Draw the Lewis electron dot diagram for each ion.
      1. Mg2+
      2. S2−
    10. Draw the Lewis electron dot diagram for each ion.
      1. In+
      2. Br−
    11. Draw the Lewis electron dot diagram for each ion.
      1. Fe2+
      2. N3−
    12. Draw the Lewis electron dot diagram for each ion.
      1. H+
      2. H−

    Answers

    1. The first two electrons in a valence shell are s electrons, which are paired.
    2.  
    3. the second column of the periodic table
    4.  
      1. 35bd44ecf0aa78e3a0893f0b805daabd.jpg
      2. 3b9dcd2bbb0c42a37149b2200aa89152.jpg
    6.  
      1. 385a984e752bb14f8267b52b63a498c5.jpg
      2. 048c529b8e83c6fc7e53a80e942fdee7.jpg
    8.  
      1. Mg2+
      2. 1176a9499078962f6f32f30092517838.jpg
    10.  
      1. Fe2+
      2. c2cba256fae0b972e562112cc9798eb9.jpg

    9.7: Electron Transfer - Ionic Bonds

    1. Comment on the possible formation of the K2+ ion. Why is its formation unlikely?
    2. Comment on the possible formation of the Cl2− ion. Why is its formation unlikely?
    3. How many electrons does a Ba atom have to lose to have a complete octet in its valence shell?
    4. How many electrons does a Pb atom have to lose to have a complete octet in its valence shell?
    5. How many electrons does an Se atom have to gain to have a complete octet in its valence shell?
    6. How many electrons does an N atom have to gain to have a complete octet in its valence shell?
    7. With arrows, illustrate the transfer of electrons to form potassium chloride from K atoms and Cl atoms.
    8. With arrows, illustrate the transfer of electrons to form magnesium sulfide from Mg atoms and S atoms.
    9. With arrows, illustrate the transfer of electrons to form scandium fluoride from Sc atoms and F atoms.
    10. With arrows, illustrate the transfer of electrons to form rubidium phosphide from Rb atoms and P atoms.
    11. Which ionic compound has the higher lattice energy—KI or MgO? Why?
    12. Which ionic compound has the higher lattice energy—KI or LiF? Why?
    1. Which ionic compound has the higher lattice energy—BaS or MgO? Why?

    Answers

    1. The K2+ ion is unlikely to form because the K+ ion already satisfies the octet rule and is rather stable.
    2.  
    3. two
    4.  
    5. two
    6.  
    7. a95eae3eeb1ce28efc326d994f6e24bf.jpg
    8.  
    9. d55643d3c011bdd321b03030380c9d1a.jpg
    10.  
    11. MgO because the ions have a higher magnitude charge
    12.  
    13. MgO because the ions are smaller

    9.8: Covalent Bonds

    1. How many electrons will be in the valence shell of H atoms when it makes a covalent bond?
    2. How many electrons will be in the valence shell of non-H atoms when they make covalent bonds?
    3. What is the Lewis electron dot diagram of I2? Circle the electrons around each atom to verify that each valence shell is filled.
    4. What is the Lewis electron dot diagram of H2S? Circle the electrons around each atom to verify that each valence shell is filled.
    5. What is the Lewis electron dot diagram of NCl3? Circle the electrons around each atom to verify that each valence shell is filled.
    6. What is the Lewis electron dot diagram of SiF4? Circle the electrons around each atom to verify that each valence shell is filled.
    7. Draw the Lewis electron dot diagram for each substance.
      1. SF2
      2. BH4−
    8. Draw the Lewis electron dot diagram for each substance.
      1. PI3
      2. OH−
    9. Draw the Lewis electron dot diagram for each substance.
      1. GeH4
      2. ClF
    10. Draw the Lewis electron dot diagram for each substance.
      1. AsF3
      2. NH4+
    11. Draw the Lewis electron dot diagram for each substance. Double or triple bonds may be needed.
      1. SiO2
      2. C2H4 (assume two central atoms)
    12. Draw the Lewis electron dot diagram for each substance. Double or triple bonds may be needed.
      1. CN−
      2. C2Cl2 (assume two central atoms)
    13. Draw the Lewis electron dot diagram for each substance. Double or triple bonds may be needed.
      1. CS2
      2. NH2CONH2 (assume that the N and C atoms are the central atoms)
    14. Draw the Lewis electron dot diagram for each substance. Double or triple bonds may be needed.
      1. POCl
      2. HCOOH (assume that the C atom and one O atom are the central atoms)

    Answers

    1. two
    2.  
    3. 9dbac5188d5eadd07ca641d5fa1d6103.jpg
    4.  
    5. e86e859ff76310d8c415dbf4cec1677b.jpg
    6.  
      1. 620716d83307accef828f56213ce474c.jpg
      2. 607d5b6c59d668c8d4085ff38a1007a6.jpg
    8.  
      1. 29bfd2b09df1cdcb26d56f57f2863746.jpg
      2. 9d7d1b8b76eaa23583aa68b909dc98a7.jpg
    10.  
      1. a27db5fad86c8228aa9d243a60e698bb.jpg
      2. 0866010fdc6513d4892e103a98de2d64.jpg
    12.  
      1. 21fd6160f13829d4b9dfc983a91b3e2d.jpg
      2. c2637a0c2491947ad6aa9426cd550c5a.jpg

    9.9: Other Aspects of Covalent Bonds

    1. Give an example of a nonpolar covalent bond. How do you know it is nonpolar?
    2. Give an example of a polar covalent bond. How do you know it is polar?
    3. How do you know which side of a polar bond has the partial negative charge? Identify the negatively charged side of each polar bond.
      1. H–Cl
      2. H–S
    4. How do you know which side of a polar bond has the partial positive charge? Identify the positively charged side of each polar bond.
      1. H–Cl
      2. N-F
    5. Label the bond between the given atoms as nonpolar covalent, slightly polar covalent, definitely polar covalent, or likely ionic.
      1. H and C
      2. C and F
      3. K and F
    6. Label the bond between the given atoms as nonpolar covalent, slightly polar covalent, definitely polar covalent, or likely ionic.
      1. S and Cl
      2. P and O
      3. Cs and O
    7. Which covalent bond is stronger—a C–C bond or a C–H bond?
    8. Which covalent bond is stronger—an O–O double bond or an N–N double bond?
    9. Estimate the enthalpy change for this reaction: N2 + 3H2 → 2NH3 .Start by drawing the Lewis electron dot diagrams for each substance.
    10. Estimate the enthalpy change for this reaction. Start by drawing the Lewis electron dot diagrams for each substance: HN=NH + 2H2 → 2NH3
    11. Estimate the enthalpy change for this reaction. Start by drawing the Lewis electron dot diagrams for each substance: CH4 + 2O2 → CO2 + 2H2O
    12. Estimate the enthalpy change for this reaction. Start by drawing the Lewis electron dot diagrams for each substance: 4NH3 + 3O2 → 2N2 + 6H2O

    Answers

    1. H–H; it is nonpolar because the two atoms have the same electronegativities (answers will vary).
    2.  
      1. Cl side
      2. S side
    4.  
      1. slightly polar covalent
      2. definitely polar covalent
      3. likely ionic
    6.  
    7. C–H bond
    8.  
    9. −80 kJ
    10.  
    11. −798 kJ

    9.10: Violations of the Octet Rule

    1. Why can an odd-electron molecule not satisfy the octet rule?
    2. Why can an atom in the second row of the periodic table not form expanded valence shell molecules?
    3. Draw an acceptable Lewis electron dot diagram for these molecules that violate the octet rule.
      1. NO2
      2. XeF4
    4. Draw an acceptable Lewis electron dot diagram for these molecules that violate the octet rule.
      1. BCl3
      2. ClO2
    5. Draw an acceptable Lewis electron dot diagram for these molecules that violate the octet rule.
      1. POF3
      2. ClF3
    6. Draw an acceptable Lewis electron dot diagram for these molecules that violate the octet rule.
      1. SF4
      2. BeH2

    Answers

    1. There is no way all electrons can be paired if there are an odd number of them.
    2.  
      1. 7652e68d149b8400c8d0a090966898fb.jpg
      2. 35c8c394c7eceb20632804aa0d42911d.jpg
    4.  
      1. 18207191d14edc34ec63b38f45bc206b.jpg
      2. 6a63a091221e0a1fcd9dc19d07212f86.jpg

    9.11: Molecular Shapes

    1. What is the basic premise behind VSEPR?
    2. What is the difference between the electron group geometry and the molecular geometry?
    3. Identify the electron group geometry and the molecular geometry of each molecule.
      1. H2S
      2. POCl3
    4. Identify the electron group geometry and the molecular geometry of each molecule.
      1. CS2
      2. H2S
    5. Identify the electron group geometry and the molecular geometry of each molecule.
      1. HCN
      2. CCl4
    6. Identify the electron group geometry and the molecular geometry of each molecule.
      1. BI3
      2. PH3
    7. What is the geometry of each species?
      1. CN−
      2. PO43−
    8. What is the geometry of each species?
      1. PO33−
      2. NO3−
    9. What is the geometry of each species?
      1. COF2
      2. C2Cl2 (both C atoms are central atoms and are bonded to each other)
    10. What is the geometry of each species?
      1. CO32−
      2. N2H4 (both N atoms are central atoms and are bonded to each other)

    Answers

    1. Electron pairs repel each other.
    2.  
      1. electron group geometry: tetrahedral; molecular geometry: bent
      2. electron group geometry: tetrahedral; molecular geometry: tetrahedral
    4.  
      1. electron group geometry: linear; molecular geometry: linear
      2. electron group geometry: tetrahedral; molecular geometry: tetrahedral
    6.  
      1. linear
      2. tetrahedral
    8.  
      1. trigonal planar
      2. linear and linear about each central atom

    9.12 Additional Exercises

    1. Explain why iron and copper have the same Lewis electron dot diagram when they have different numbers of electrons.
    2. Name two ions with the same Lewis electron dot diagram as the Cl− ion.
    3. Based on the known trends, what ionic compound from the first column of the periodic table and the next-to-last column of the periodic table should have the highest lattice energy?
    4. Based on the known trends, what ionic compound from the first column of the periodic table and the next-to-last column of the periodic table should have the lowest lattice energy?
    5. P2 is not a stable form of phosphorus, but if it were, what would be its likely Lewis electron dot diagram?
    6. Se2 is not a stable form of selenium, but if it were, what would be its likely Lewis electron dot diagram?
    7. What are the Lewis electron dot diagrams of SO2, SO3, and SO42−?
    8. What are the Lewis electron dot diagrams of PO33− and PO43−?
    9. Which bond do you expect to be more polar—an O–H bond or an N–H bond?
    10. Which bond do you expect to be more polar—an O–F bond or an S–O bond?
    11. Use bond energies to estimate the energy change of this reaction. C3H8 + 5O2 → 3CO2 + 4H2O
    12. Use bond energies to estimate the energy change of this reaction. N2H4 + O2 → N2 + 2H2O
    13. Ethylene (C2H4) has two central atoms. Determine the geometry around each central atom and the shape of the overall molecule.
    14. Hydrogen peroxide (H2O2) has two central atoms. Determine the geometry around each central atom and the shape of the overall molecule.

    Answers

    1. Iron has d electrons that typically are not shown on Lewis electron dot diagrams.
    2.  

    3. LiF

    4.  

    5. It would be like N2:

    6.  
    8.  

    9. an O–H bond

    10.  

    11. −2,000 kJ

    12.  

    13. trigonal planar about both central C atoms

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