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Atomic Structure & Bonding AP Sample Problems
1.
The effective nuclear charge experienced by the outermost electron of Na is different
than the effective nuclear charge experienced by the outermost electron of Ne. This
difference best accounts for which of the following? (0246)
(A)
(B)
(C)
(D)
(E)
Na has a greater density at standard conditions than Ne.
Na has a lower first ionization energy than Ne.
Na has a higher melting point than Ne.
Na has a higher neutron-to-proton ratio than Ne.
Na has fewer naturally occurring isotopes than Ne.
Questions 2-5 refer to atoms for which the occupied atomic orbitals shown below. (995-8)
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2. Represents an atom that is chemically unreactive
3. Represents an atom in an excited state
4. Represents an atom that has four valence electrons.
5. Represents an atom of a transition metal.
6. In the periodic table, as the atomic number increases from 11 to 17, what happens to the
atomic radius? (9950)
(A) It remains constant.
(B) It increases only.
(C) It increases, then decreases.
(D) It decreases only.
(E) It decreases, then increases.
Questions 7-10
(941-4)
(A) Heisenberg uncertainty principle
(B) Pauli exclusion principle
(C) Hund's rule (principle of maximum multiplicity)
(D) Shielding effect
(E) Wave nature of matter
7. Can be used to predict that a gaseous carbon atom in its ground state is paramagnetic
8. Explains the experimental phenomenon of electron diffraction
9. Indicates that an atomic orbital can hold no more than two electrons
10. Predicts that it is impossible to determine simultaneously the exact position and the exact
velocity of an electron
11. In a molecule in which the central atom exhibits sp3d2 hybrid orbitals, the electron pairs are
directed toward the corners of
(9415)
(A) a tetrahedron
(B) a square-based pyramid
(C) a trigonal bipyramid
(D) a square
(E) an octahedron
12. All of the following statements concerning the characteristics of the halogens are true
EXCEPT: (9454)
(A) The first ionization energies (potentials) decrease as the atomic numbers of the
halogens increase.
(B) Fluorine is the best oxidizing agent.
(C) Fluorine atoms have the smallest radii.
(D) Iodine liberates free bromine from a solution of bromide ion.
(E) Fluorine is the most electronegative of the halogens.
13. Molecules that have planar configurations include which of the following? (9457)
I. BCl3
II. CHCl3
III. NCl3
(A) I only
(B) III only
(C) I and II only
(D) II and III only
(E) I, II, and III
Use these answers for questions 14 - 16. (891-3)
(A) O
(B) La
(C) Rb
(D) Mg
(E) N
14. What is the most electronegative element of the above?
15. Which element exhibits the greatest number of different oxidation states?
16. Which of the elements above has the smallest ionic radius for its most commonly found ion?
17. The Lewis dot structure of which of the following molecules shows only one unshared pair
of valence electron? (8917)
(A) Cl2
(B) N2
(C) NH3
(D) CCl4
(E) H2O2
18. CCl4, CO2, PCl3, PCl5, SF6 Which of the following does not describe any of the molecules
above? (8947)
(A) Linear
(B) Octahedral
(C) Square planar
(D) Tetrahedral
(E) Trigonal pyramidal
19. The geometry of the SO3 molecule is best described as (8440)
(A) trigonal planar
(B) trigonal pyramidal
(C) square pyramidal
(D) bent
(E) tetrahedral
20. Which of the following molecules has the shortest bond length? (8441)
(A) N2
(B) O2
(C) Cl2
(D) Br2
(E) I2
21. The elements in which of the following have most nearly the same atomic radius? (8443)
(A) Be, B, C, N
(B) Ne, Ar, Kr, Xe
(C) Mg, Ca, Sr, Ba
(D) C, P, Se, I
(E) Cr, Mn, Fe, Co
22. Pi bonding occurs in each of the following species EXCEPT
(8451)
(A) CO2
(B) C2H4
(C) CN¯
(D) C6H6
(E) CH4
23. Which of the following represents the ground state electron configuration for the Mn3+ ion?
(Atomic number Mn = 25)
(8458)
(A) 1s2 2s22p6 3s23p63d4
(B) 1s2 2s22p6 3s23p63d5 4s2
(C) 1s2 2s22p6 3s23p63d2 4s2
(D) 1s2 2s22p6 3s23p63d8 4s2
(E) 1s2 2s22p6 3s23p63d3 4s1
24. Which of the following has a zero dipole moment? (8460)
(A) HCN
(B) NH3
(C) SO2
(D) NO2
(E) PF5
Free Response Samples
1) The values of the first three ionization energies (I1, I2, I3) for magnesium and argon [in
kJ/mole] are as follows:
I1
I2
I3
Mg
735
1443
7730
Ar
1525
2665
3945
(a) Give the electronic configuration of Mg and Ar.
(b) In terms of these configurations, explain why the values of the first and second
ionization energies of Mg are significantly lower than the values for Ar, whereas the third
ionization energy of Mg is much larger than the third ionization energy Ar.
(c) If a sample of Ar in one container and a sample of Mg in another container are each
heated and chlorine is passed in to each container, what compounds, if any , will be
formed? Explain in terms of the electronic configuration given in part (a).
(d) Element Q has the following first three ionization energies [in kJ/mole]:
I1
I2
I3
Q
496
4568
6920
What is the formula for the most likely compound of element Q with chlorine? Explain the
choice of formula on the basis of the ionization energies. (82-6)
2) Two important concepts that relate to the behavior of electrons in atomic system are the
Heisenberg uncertainty principle and the wave-particle duality of matter. (87-9)
a) State the Heisenberg uncertainty principle as it relates to determining the position and
momentum of an object.
b) What aspect of the Bohr theory of the atom is considered unsatisfactory as a result of
the Heisenberg uncertainty principle?
c) Explain why the uncertainty principle or the wave nature of particles is not significant
when describing the behavior of macroscopic objects, but is very significant when
describing the behavior of electrons.
3)
CF4
XeF4
ClF3
(a) Draw a Lewis electron-dot structure for each of the molecules above and identify the shape of
each. (89-5)
(b) Use the valence shell electron-pair repulsion (VSEPR) model to explain the geometry of each
of these molecules.
4)
Element
lithium
beryllium
boron
carbon
nitrogen
oxygen
fluorine
neon
First Ionization Energy (kJ/mol)
520
899
800
1086
1402
1314
1681
2088
The diagram shows the first ionization energies for the elements from Li to Ne. Briefly (in one to
three sentences) explain each of the following in terms of atomic structure. (90-6)
(a) In general, there is an increase in the first ionization energy from Li to Ne.
(b) The first ionization energy of B is lower than that of Be.
(c) The first ionization energy of O is lower than that of N.
(d) Predict how the first ionization energy of Na compares to those of Li and of Ne.
Explain.
5)
NO2
NO2¯
NO2+
Nitrogen is the central atom in each of the species given above. (92-9)
(a) Draw the Lewis electron-dot structure for each of the three species.
(b) List the species in order of increasing bond angle. Justify your answer.
(c) Select one of the species and give the hybridization of the nitrogen atom in it.
(d) Identify the only one of the species that dimerizes and explain what causes it to do so.
6) Account for each of the following in terms of principles of atomic structure, including the
number, properties, and arrangements of subatomic particles. (93-6)
(a) The second ionization energy of sodium is about three time greater than the second ionization
energy of magnesium.
(b) The difference between the atomic radii of Na and K is relatively large compared to the
difference between the atomic radii of Rb and Cs.
(c) A sample of solid nickel chloride is attracted into a magnetic field, whereas a sample of solid
zinc chloride is not.
(d) Phosphorus forms the fluorides PF3 and PF5, whereas nitrogen forms only NF3.
7) Use principle of atomic structure and/or chemical bonding to answer of each of the following.
(94-9)
(a) The radius of the Ca atom is 0.197 nanometer; the radius of the Ca2+ ion is 0.099 nanometer.
Account for this difference.
(b) The lattice energy of CaO(s) is -3,460 kilojoules per mole; the lattice energy for K2O(s) is 2,240 kilojoules per mole. Account for this difference.
Ionization Energy
First
Second
K
419
3,050
Ca
590
1,140
(c) Explain the difference between Ca and K in regard to:
(i) their first ionization energies.
(ii) their second ionization energies.
(d) The first ionization energy of Mg is 738 kilojoules per mole and that of Al is 578 kilojoules
per mole. Account for this difference.
8) Explain each of the following obsevations using principles of atomic structure and/or bonding.
(97-6)
a) Potassium has a lower first-ionization energy than lithium.
b) The ionic radius of N3¯ is larger than that of O2¯.
c) A calcium atom is larger than a zinc atom.
d) Boron has a lower first-ionization energy than beryllium.
Answer the following questions using principles of chemical bonding and molecular structure.
(99-8ef)
9) a) Consider the carbon dioxide molecule, CO2 , and the carbonate ion, CO32-.
i. Draw the complete Lewis electron-dot structure for each species.
ii. Account for the fact that the carbon-oxygen bond length in CO32- is greater than
the carbon-oxygen bond length in CO2.
b) Consider the molecules CF4 and SF4.
iii. Draw the complete Lewis electron-dot structure for each molecule.
iv. In terms of molecular geometry, account for the fact that the CF4 molecule is
nonpolar, whereas the SF4 molecule is polar.
10)
Answer the following questions about the element selenium, Se (atomic number 34).
(00-07)
a)
Samples of natural selenium contain six stable isotopes. In terms of atomic
structure, explain what these isotopes have in common, and how they differ.
b)
Write the complete electron configuration (e.g., 1s2 2s2 … etc.) for a selenium
atom in the ground state. Indicate the number of unpaired electrons in the ground-state atom,
and explain your reasoning.
c)
In terms of atomic structure, explain why the first ionization energy of selenium is
i)
less than that of bromine (atomic number 35), and
ii)
greater than that of tellurium (atomic number 52).
d)
Selenium reacts with fluorine to form SeF4. Draw the complete Lewis electrondot structure for SeF4 and sketch the molecular structure. Indicate whether the molecule is polar
or nonpolar, and justify your answer.
11. Use the principles of atomic structure and/or chemical bonding to explain each of the
following. In each part, you answer must include references to both substances. (02-06)
a.) The atomic radius of Li is larger than that of Be.
b.) The second ionization energy of K is greater than the second ionization energy of
Ca.
c) The carbon-carbon bond energy in C2H4 is greater than it is in C2H6.
12.
Answer the following questions that relate to chemical bonding. (05-06)
a)
In the boxes provided, draw the complete Lewis structure (electron-dot diagram)
for each of the three molecules represented below.
CF4
b)
c)
PF5
SF4
On the basis of the Lewis structures drawn above, answer the following questions
about the particular molecule indicated.
i)
What is the F-C-F bond angle in CF4?
ii)
What is the hybridization of the valence orbitals of P in PF5?
iii)
What is the geometric shape formed by the atoms in SF4?
Two Lewis structures can be drawn for the OPF3 molecule, as shown below.
i)
How many sigma bonds and how many pi bonds are in structure 1?
ii)
Which one of the two structures best represents a molecule of OPF3?
Justify your answer in terms of formal charge.
13.
Use principles of atomic structure, bonding, and/or intermolecular forces to respond to
each of the following. Your responses must include specific information about all substances
referred to in each question. (07-c
c)
As shown in the table below, the first ionization energies of Si, P, and Cl show a trend.
Element
First Ionization Energy
(kJ mol-1)
Si
786
P
1,012
Cl
1,251
i)
For each of the three elements, identify the quantum level (e.g. n = 1, n =
2, etc.) of the valence electrons in the atom.
ii)
Explain the reasons for the trend in first ionization energies