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  • 34.

    Discuss the dual wave-particle nature of light.

  • 35.

    Discuss the significance of the photoelectric effect and the line-emission spectrum of hydrogen to the development of the atomic model.

  • 36.

    Describe the Bohr model of the hydrogen atom.

  • 37.

    Discuss Louis de Broglie’s role in the development of the quantum model of the atom.

  • 38.

    Compare and contrast the Bohr model and the quantum model of the atom.

  • 39.

    Explain how the Heisenberg uncertainty principle and the Schrödinger wave equation led to the idea of atomic orbitals.

  • 40.

    List the four quantum numbers, and describe their significance.

  • 41.

    Relate the number of sublevels corresponding to each of an atom’s main energy levels, the number of orbitals per sublevel, and the number of orbitals per main energy level.

  • 42.

    List the total number of electrons needed to fully occupy each main energy level.

  • 43.

    State the Aufbau principle, the Pauli exclusion principle, and Hund’s rule.

  • 44.

    Describe the electron configurations for the atoms of any element using orbital notation, electron- configuration notation, and, when appropriate, noble-gas notation.

Chapter 5 The Periodic Law

  • 45.

    Explain the roles of Mendeleev and Moseley in the development of the periodic table.

  • 46.

    Describe the modern periodic table.

  • 47.

    Explain how the periodic law can be used to predict the physical and chemical properties of elements.

  • 48.

    Describe how the elements belonging to a group of the periodic table are interrelated in terms of atomic number.

  • 49.

    Describe the relationship between electrons in sublevels and the length of each period of the periodic table.

  • 50.

    Locate and name the four blocks of the periodic table. Explain the reasons for these names.

  • 51.

    Discuss the relationship between group configurations and group numbers.

  • 52.

    Describe the locations in the periodic table and the general properties of the alkali metals, the alkaline- earth metals, the halogens, and the noble gases.

  • 53.

    Define atomic and ionic radii, ionization energy, electron affinity, and electronegativity.

  • 54.

    Compare the periodic trends of atomic radii, ionization energy, and electronegativity, and state the reasons for these variations.

  • 55.

    Define valence electrons, and state how many are present in atoms of each main-group element.

  • 56.

    Compare the atomic radii, ionization energies, and electronegativities of the d-block elements with those of the main-group elements.

Chapter 6: Chemical Bonding

  • 57.

    Define chemical bond.

  • 58.

    Explain why most atoms form chemical bonds.

  • 59.

    Describe ionic and covalent bonding.

  • 60.

    Explain why most chemical bonding is neither purely ionic nor purely covalent.

  • 61.

    Classify bonding type according to electronegativity differences.

  • 62.

    Define molecule and molecular formula.

  • 63.

    Explain the relationships between potential energy, distance between approaching atoms, bond length, and

bond energy.

  • 64.

    State the octet rule.

  • 65.

    List the six basic steps used in writing Lewis structures.

  • 66.

    Explain how to determine Lewis structures for molecules containing single bonds, multiple bonds, or both.

  • 67.

    Explain why scientists use resonance structures to represent some molecules.

  • 68.

    Compare and contrast a chemical formula for a molecular compound with one for an ionic compound.

  • 69.

    Discuss the arrangements of ions in crystals.

  • 70.

    Define lattice energy and explain its significance.

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