7. What is an electron configuration? Give an example.
An electron configuration shows the particular orbitals that are occupied by electrons in an atom.
Cu-> [Ar] 4s^23d^9
8. What is columb's law? Explain how the potential energy of two charged particles depends on the distance between there charged particles and on the magnitude and sign of their charges.
Describes the attractions and repulsions between charged particles. For like charges, the potential energy is positive and decreases as the particles get further apart. For opposite charges, the potential energy is negative and becomes more negative as th
9. What is shielding? In an atom, which electrons tend to do the most shielding (core or valence?)
Repulsions in multi-electron atoms that cause the electron to have a net reduced attraction to the nucleus.
Shielding or screening occurs when one electron is blocked from the full effects of the nuclear charge so that the electron experiences only a part
10. What is penetration? How does the penetration of an orbital into the region occupied by core electrons affect the energy of an electron in that orbital?
Penetration occurs when an electron penetrates the electron cloud of the the 1s orbital and experiences the charge of the nucleus more fully because it is less shielded by the intervening electrons.
As the outer electron undergoes penetration into the reg
11. Why are the sub levels within a principal level split into different energies for multi-electron atoms but not for the hydrogen atom?
The sublevels within a principle level split in multielectron atoms because of penetration of the outer electrons into the region of the core electrons.
The sublevels in hydrogen are not split because they are empty in the ground state.
12. What is an orbital diagram?
An orbital diagram is a different way to show the electron configuration of an atom. It symbolizes the electron as an arrow in a box that represents the orbital.
The orbital for a hydrogen atom: (is a box with 1s under it and an H next to it. It has one a
13. Why is electron spin important when writing electron configurations? Explain in terms of the Pauli exclusion principle.
The Pauli exclusion principle states the following: "No two electrons in an atom can have the same four quantum numbers."
Because two electrons occupying the same orbital have three identical quantum numbers (n,l,m) they must have different spin quantum n
14. What are degenerate orbitals? According to Hund's rule, how are degenerate orbital occupied?
Degenerate orbitals are orbitals of the same energy. In a multi-electron atom, the orbitals in a sub-levels are degenerate.
Hund's rule states that when filling degenerate orbitals, electrons fill them singly first , with parallel spins. This is the resul
15. List all orbitals from 1s through 5s according to increasing energy for multi-election atoms.
1s2s2p3s3p4s3d4p5s4d5p6s4f5d6p
16. What are valence electrons? Why are they important?
Valence electrons are important in chemical bonding. For main-group elements, the valence electrons are in the outermost principle energy level. For transition elements, we also count the outermost d electrons among the valence even though they are not in
18. Explain why the s block in the periodic table has only two columns while the p block has six.
The number of columns in a block corresponds to the maximum number of electrons that can occupy the particular sublevels of that block
. The s block has two columns corresponding to one of the s orbitals holding a maximum of two electrons.
The p block has
24. Describe the relationship between the properties of an element and the number of valence electrons that it contains.
The chemical properties of elements are largely determined by the number of valence electrons they contain. Their properties are periodic because the number of valence electrons is periodic.
Because elements within a column in the periodic table have the
28. Use the concepts of effective nuclear charge, shielding, and n value of the valence orbital to explain the trend in atomic radius as you move across a period in the periodic table.
As you move to the right across a row in the periodic table, the n level stays the same. However, the nuclear charge increases and the amount of shielding stays about the same because the number of inner electrons stays the same.
So the effective nuclear
29. For transition elements, describe and explain the observed trends in atomic radius as you move:
a. across a period in the periodic table
b. down a column in the period table
a. Atoms decrease in size from left to right across a period
b. Atoms in the same group increase in size down the column
30. How is the electron configuration of an anion different from that of the corresponding neutral atom? How is the electron configuration go a cation different?
The electron configuration of a main-group monatomic ion can be deduced from the electron configuration of the neutral atom and the charge of the ion.
For anions, we simply add the number of electrons required by the magnitude of the charge of the anion.
32. Describe the relationship between
a. the radius of a cation and that of the atom from which it forms
b. the radius of an anion and that of the atom from which it forms
a. cations are much smaller than their corresponding parent (they lose electrons from valence shell)
b. anions are much larger than their corresponding parent (they gain electrons)
34. What is the general trend in first ionization energy as you move down a column in the periodic table? As you move across a row?
As you moves down a column the ionization energy decreases. As you move across a row the ionization energy increases.
35. What are the exceptions to the periodic trends in the first ionization energy? Why do they occur?
Be,B :
To ionize Be you must break up a full sublevel, costs extra energy. When you ionize B you get a full sublevel, costs less energy
N,O:
To ionize N you must break up a half-full sublevel, costs extra energy. When you ionize O you get a half-full subl
36. Examination of the first few successive ionization energies for a given element usually reveals a large jump between two ionization energies. For example, the successive ionization energies of magnesium show a large jump between IE2 and IE3. The succe
First ionization energy of sodium involves removing valence electrons > First ionization energy is very low. Second ionization energy involves removing a core electron from an ion with a noble gas configuration > this requires a lot of energy making IE2 v
37. What is electron affinity? What are the observed periodic trends in electron affinity?
The energy change associated with the gaining of an electron by the atom in the gaseous state.The energy change associated with the gaining of an electron by the atom in the gaseous state.
As we move down a column it varies but as we move across a row the
39. Write a general equation for the reaction of an alkali metal with each substance.
a. Halogen
b. Water
a. K(s) + Br2(g) -->
K(s) + Br2(g) -->K+ Br-
2 K(s) + Br2(g) -->2 KBr(s)
( Alkali metals are oxidized to the 1+ ion. Halogens are reduced to the 1- ion. The ions then attach together by ionic bonds. The reaction is exothermic.)
b. Rb(s) + H2O(l)
Rb(s) + H
43. Write a full orbital diagram for each element.
a. N
b. F
c. Mg
d. Al
a. 1s^2 2s^2 2p^3 (p boxes only half filled)
b. [He] 2s^2 2p^5 (p boxes missing only 1)
c. [Ne] 3s^2 (3s boxes full)
d. [Ne] 3s^2 3p^1 (3p box missing 5)
46. Use the periodic table to determine the element corresponding to each electron configuration.
a. [Ar] 4s^2 3d^10 4p^6
b. [Ar] 4s^2 3d^3
c. [Kr] 5s^2 4d^10 5p^2
d. [Kr] 5s^2
a. Kr
b. V
c. Sn
d. Sr
48. Use the periodic table to determine each quantity.
a. the number of 3s electrons in Mg
b. the number of 3d electrons in Cr
c. the number of 4d electrons in Y
d. the number of 6p electrons in Pb
a. 2
b. 4
c. 1
d. 2
53. Which outer electron configuration would you expect to belong to a reactive metal? To a reactive nonmetal?
a. ns^2
b. ns^2 np^6
c. ns^2 np^5
d.ns^2 np^2
Metal: ns^2 and ns^2 np^2
Nonmetal: ns^2 np^5
55. According to Coulumb's Law, which pair of charged particle has the lowest potential energy?
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56. Acccording to Coulumb's Law, rank the interactions between charged particles from lowest potential energy to highest potential energy.
a. a 1+ charge and a 1- charge separated by 100 pm
b. a 2+ charge and a 1- charge separ, v gated by 100 pm
c. a 1+ c
Lowest -> Highest
2+ charge and 1- charge separated by 100 pm (-.0002)
1+ charge and 1- charge separated by 100 pm (-.0001)
1+ charge and 1- charge separated by 200 pm (-0.000025)
1+ charge and 1+ charge separated by 100 pm (0.0001)
(charge x charge)/(pm^
57. Which experience a greater effective nuclear charge: the valence electrons in beryllium or the valence electrons in nitrogen? Why?
The valence electrons in nitrogen as Zeff= 7-3=5+ ( total - core = valence) compared to beryllium Zeff= 4-2= 2+
59. If core electrons completely shielded valence electrons from nuclear charge (i.e., if each core electron reduced nuclear charge by 1 unit) and if valence electrons didn't not shield one another from nuclear charge at all, what would be the effective n
a. One
b. Three
c. Six
d. Four
61. Choose the larger atom from each pair.
a. Al or In
b. Si or N
c. P or Pb
d. C or F
a. In
b. Si
c. Pb
d. C
Increase ->
Decrease Down
62. Choose the larger atom from each pair, if possible.
a. Sn or Si
b. Br or Ga
c. Sn or Bi
d. Se or Sn
a. Sn
b. Ga
c. Bi ( though opposing trends, the increase down in size is greater than the decrease of moving to the right)
d. Sn
63. Arrange these elements in order of increasing atomic radius: Ca, Rb, S, Si, Ge, F.
Rb > Ca > Ge > Si > S > F
64. Arrange these elements in order of decreasing atomic radius: Cs, Sb, S, Pb, Se.
Cs > Pb > Sb > Se > S
65. Write the electron configuration for each ion.
a. O^2-
b. Br^-
c. Sr^2+
d. Cu^2+
a. 1s^2 2s^2 2p^6
b. Kr or 1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^10 4p^6
c. Kr
d. [Ar] 4s^2 3d^7
67. Write orbital diagrams for each ion and indicate whether the ion is diamagnetic or paramagnetic.
a. V^5+
b. Cr^3+
c. Ni^2+
d. Fe^3+
a. Diamagnetic
b. Diamagnetic
c. Paramagnetic
d. Paramagnetic
68. Write orbital diagrams for each ion and indicate whether the ion is diamagnetic or paramagnetic.
a. Cd^2+
b. Au^+
c. Mo^3+
d. Zr^2+
a. Paramagnetic
b. Paramagnetic
c. Exception- Paramagnetic
d. Diamagnetic
69. Which is the larger species in each pair?
a. Sr or Sr^2+
b. N or N^3-
c. Ni or Ni^2+
d. S^2- or Ca^2+
a. Sr
b. N
c. Ni
d. S^2-
71. Arrange this isoelectronic series in order of decreasing radius: F^-, O^2-, Mg^2+, Na^+.
O^2-, F^-, Na^+, Mg^2+
73. Choose the element with the higher first ionization energy from each pair.
a. Br or Bi
b. Na or Rb
c. As or At
d. P or Sn
Increasing up and right
a. Br
b. Na
c. opposing trends
d. P
76. Arrange these elements in order of decreasing first ionization energy: Cl, S, Sn, Pb.
Cl -> S -> Sn -> Pb
80. Choose the element with the more negative (more exothermic) electron affinity from each pair.
a. Mg or S
b. K or Cs
c. Si or P
d. Ga or Br
a.
b.
c.
d.
81. Choose the more metallic element from each pair.
a. Sr or Sb
b. As or Bi
c. Cl or O
d. S or As
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83. Arrange these elements in order of increasing metallic character: Fr, Sb, In, S, Ba, Se.
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91. Bromine is highly reactive liquid while krypton in an inert gas. Explain this different based on their electron configurations.
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93. Both vanadium and its 3+ ion are paramagnetic. Refer to their electron configurations to explain this statement.
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95. Suppose you were trying to find a substitute for K+ in nerve signal transmission. Where would you being your search? What ions would be most like K+? For each ion you propose, explain the ways in which it would be similar to K+ and the ways it would b
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98. Which pair of elements would you expect to have the most similar atomic radii, and why?
a. Si and Ga
b. Si and Ge
c. Si and As
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99. Consider these elements: N, Mg, O, F, Al.
a. Write the electron configuration for each element.
b. Arrange the elements in order of decreasing atomic radius.
c. Arrange the elements in order of increasing ionization energy.
d. Use the electron configu
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100. Consider these elements: P, Ca, Si, S, Ga.
a. Write the electron configuration for each element.
b. Arrange the elements in order of decreasing atomic radius.
c. Arrange the elements in order of increasing ionization energy.
d. Use the electron confi
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101. Explain why atomic radius decreases as you move to the right across a period for main-group elements but not for tradition elements.
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102. Explain why vanadium (radius = 134 pm) and copper (radius = 128 pm) have nearly identical atomic radii, even though the atomic number of copper is about 25% higher than that of vanadium. What would you predict about the relative densities of these tw
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104. The lightest halogen is als o the most chemically reactive, and reactivity generally decreases as you move down the column of halogens in the periodic table. Explain this trend in terms of periodic properties.
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105. Write general outer electron configurations for groups 6a and 7a in the periodic table. The electron affinity of each group 7a element is more negative than that of each corresponding group 6a element. Use the electron configurations to explain why t
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110. The electron affinity of sodium is lower than that of lithium, while the electron affinity of chlorine is higher than that of fluorine. Suggest an explanation for this observation.
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111. Use Columb's law to calculate the ionization energy in kJ/mol of an atom composed of a proton and an electron separated by 100pm. what wavelength of light has sufficient energy to ionize the atom?
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113. Consider the elements: B, C, N, O, F.
a. which element has the highest first ionization energy?
b. Which element has the largest atomic radius?
c. Which element is most metallic?
d. Which element has three unpaired electrons?
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119. A carbon atom can absorb radiation of various wavelengths with resulting changes in its electron configuration. Write orbital diagrams for the electron configuration of carbon that results from absorption of the three longest wavelengths of radiation
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