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CH185, General Chemistry I

Exam #3, March 29, 1996

Short Answer (Problems 1-14 worth 4 points each)

1. Write the ground electron configuration of an Si atom      [Ne]3s²3p²

2. Write the ground electron configuration of the Co³⁺ ion      [Ar]3d⁶

3. How many unpaired electrons does an atom of S have?   2   ([Ne]3s²3p⁴ configuration)

4. Write the following ions in order of increasing (from smallest to largest) ionic radius: O^2-, Se^2-, F^-, S^2-,Te^2-
F^-, O^2-, S^2-, Se^2-, Te^2-
5. What values of the angular momentum quantum number (l) are allowed for a principal quantum number (n) of 4?
0, 1, 2, and 3
6. In each of the following pairs of elements, circle the chemical symbol of the element that has the higher ionization energy.   (Correct answers are underlined here with comment on line below)

     (a) Ne or Ar         (b) Be or B          (c) C or N           (d) S or P 
    (Ne above Ar)      (exception to trend,  (N to right of C)   (exception to trend, 
                     B has e^- in p subshell)                    P has exactly 1/2-filled p s.s)

7. Are the Ga-Cl bonds in the solid compound GaCl₃ ionic, polar covalent, or nonpolar covalent bonds? (Electronegativity of Ga = 1.6; electronegativity of Cl = 3.0)

(E.N.) = 3.0 - 1.6 = 1.4 < 2.0, so bond is polar covalent
8. Use Lewis dot symbols for the reactants and products in a chemical equation to show the transfer of electrons when the elements Al and F react to form the ionic solid compound AlF₃.

Write Lewis dot structures in which the octet rule is obeyed for each of the following molecules or polyatomic ions in problems 9-12.
9. O₂ 2(6) = 12 e^-

10. BrO₂^- 7 + 2(6) + 1 = 20 e^-

O
11. HCOOH -- skeletal structure: H C O H 1 + 4 + 12 + 1 = 18 e^-

12. BF₄^- 3 + 4(7) + 1 = 32 e^-

13. For the given Lewis dot structure, write the oxidation number of each nitrogen atom on the line below the atom. Electronegativity of N = 3.0; of H = 2.1. Use your knowledge of trends in electronegativity to determine its relative value among the elements C, N, and O

Assignment of bonding electrons to more electronegative atoms or split between atoms of equal electronegativity is shown on the diagram below.

oxidation number = (# of valence electrons) - (# of electrons assigned in Lewis structure)
oxidation number: 5-8 = -3 5-2 = +3

14. Write the formal charge next to each atom in the following two Lewis dot structures of SO₄^2-, then circle the Lewis dot structure that is more plausible on the basis of formal charge.

Atomic Orbitals
(8 points) 15. For each of the boundary surface diagrams of atomic orbitals shown below, write the small letter designation for the orbital and the number of orbitals allowed in a particular sublevel of that orbital type.

small letter designation:       d             p             d               s

number of orbitals 
in subshell:                    5             3              5               1

Word Problems (Point value listed to the left of each problem number).
(8 points) 16. Phosphorescent molecules emit light (``glow in the dark'') when an electron jumps from a higher energy orbital to a lower energy orbital. If a particular phosphorescent molecule emits yellow light with a wavelength of 555 nm, what is the difference in energy between these two orbitals? (c = 3.00 x 10⁸ m/s; h = 6.63 x 10^-34 J-s; 1 nm = 1 x 10^-9 m).

E = E(photon) = h

= hc/

= [(6.63 x 10^-34 J-s)(3.00 x 10⁸ m/s)]/(555 x 10^-9 m) = 3.58 x 10^-19 J

(13 points) 17. The Born-Haber cycle is a means of finding the lattice energy of ionic solids through application of Hess' Law to a series of steps that lead from the ionic solid to the gaseous ions. Calculate the lattice energy of RbCl(s) given the following thermochemical equations.

Overall chemical equation: RbCl(s) --> Rb⁺(g) + Cl^-(g)

H = ??????

Rb(s) + 1/2Cl₂(g) --> RbCl(s)

H = -430.6 kJ <--reverse
Rb(s) --> Rb(g)

H = 85.8 kJ
Cl₂(g) --> 2Cl(g)

H = 242.0 kJ <--divide by 2
Rb(g) --> Rb⁺(g)

H = 409.2 kJ
Cl(g) --> Cl^-(g)

H= -349.0 kJ

RbCl(s) --> Rb(s) + 1/2Cl₂(g)

H = +430.6 kJ
Rb(s) --> Rb(g)

H = 85.8 kJ
1/2Cl₂(g) --> Cl(g)

H = 121.0 kJ
Rb(g) --> Rb⁺(g)

H = 409.2 kJ
Cl(g) --> Cl-(g)

H= -349.0 kJ
RbCl(s) --> Rb⁺(g) + Cl^-(g)

H =   697.6 kJ

Short Essay (5 points each) Please answer each question with a 2-3 sentence explanation.
18. What general properties of an atomic orbital do the principal quantum number (n), the angular momentum quantum number (l), and the magnetic quantum number (m_l) each describe?

1. Principal quantum number decribes size and approximate energy of orbital.

2. Angular momentum quantum number describes shape of orbital.

3. Magnetic quantum number describes orientation of orbital in space.

19. Why do the atomic radii of the elements decrease as one moves from left to right across a row of the periodic table in spite of the fact that the number of protons and electrons is increasing?

     Electrons are added only to the valence orbitals where they do not shield the nuclear charge effectively. Hence elements further to the right in a period have larger effective nuclear charges that exert stronger attractive forces on the valence electrons. These stronger attractive forces pull the electrons closer to the nucleus, resulting in a smaller radius for the atom.

20. Can hydrogen atoms in a Lewis dot structure have lone pair electrons or multiple bonds? Explain why or why not.

      Hydrogen fills its valence shell (1s) and achieves a noble gas configuration with just two electrons. This is accomplished when hydrogen forms a single covalent bond. The presence of lone pair electrons or multiple bonds would give the hydrogen atom more electrons than needed to fill its valence shell. Hence hydrogen never has lone pair electrons or participates in multiple bonds.