Topic 6 ELECTRON CONFIGURATION practice problems
1. The atom having the valence-shell configuration 4s2 4p5 would be in:
(a) Group 17 and Period 5
(b) Group 14 and Period 4
(c) Group 6 and Period 7
(d) Group 7 and Period 4
(e) Group 17 and Period 4 Br has configuration of 1s22s22p63s23p64s23d104 p5
2. Select the term best describing the series of elements: Mn, Fe, Co, Ni, Cu.
(a) d-transition metals
(b) representative elements
(c) metalloids
(d) alkaline earth metals
(e) halogens
3. The species that contains 24 protons, 26 neutrons and 22 electrons would be represented by
the symbol:
(a) 50V3+
(b) 26Cr2+
(c) 50Cr2+
(d) 50Mn2+
(e) none of these
4. 1s2 2s22p6 3s23p3
Atoms of an element, X, have the electronic configuration shown above. The compound most
likely formed with magnesium, Mg, is…
(A) MgX
(B) Mg2X
(C) MgX2 (D) MgX3 (E) Mg3X2
5. Which of the following represents the ground state electron configuration for the Mn 3+ ion?
(Atomic number Mn = 25)
(A) 1s2 2s22p6 3s23p63d4
(B) 1s2 2s22p6 3s23p63d5 4s2 (C) 1s2 2s22p6 3s23p63d2 4s2
(D) 1s2 2s22p6 3s23p63d8 4s2 (E) 1s2 2s22p6 3s23p63d3 4s1
Use these answers for questions 6 - 9.
(A) 1s2 2s22p5 3s23p5
(B) 1s2 2s22p6 3s23p6
(C) 1s2 2s22p62d10 3s23p6
2
2
6
2
6
5
2
2
6
2
6
3
2
(D) 1s 2s 2p 3s 3p 3d
(E) 1s 2s 2p 3s 3p 3d 4s
6. An impossible electronic configuration C (no 2d)
7. The ground-state configuration for the atoms of a transition element E
8. The ground-state configuration of a negative ion of a halogen B
9. The ground-state configuration of a common ion of an alkaline earth element B
Use these answers for questions 10-14
(A) Heisenberg uncertainty principle (B) Pauli exclusion principle
(C) Hund's rule
(D) Shielding effect
(E) Wave nature of matter (F) Aufbau Principle
10. Can be used to predict that a gaseous carbon atom in its ground state contains an unpaired
electron C
11.. Explains the experimental phenomenon of electron diffraction E
12. Indicates that an atomic orbital can hold no more than two electrons B
13. Predicts that it is impossible to determine simultaneously the exact position and the exact
velocity of an electron A
14. As electrons are added to an atom, they will assume the most stable condition and fill in
lower energy levels before filling higher energy levels F
Questions 15-18 refer to atoms for which the occupied atomic orbitals are shown below:
15. Represents an atom that is chemically unreactive D
16. Represents an atom in an excited state A
17. Represents an atom that has four valence electrons. C
18. Represents an atom of a transition metal. E
Consider atoms of the following elements for questions 19-20. Assume that the atoms are in
the ground state.
(A) S
(B) Ca
(C) Ga
(D) Sb
(E) Br
19. The atom that contains exactly two unpaired electrons. A
20. The atom that contains only one electron in the highest occupied energy sublevel. C
21. In which of the following groups are the three species isoelectronic?
(A) S2¯, K+, Ca2+
(B) Sc, Ti, V2+
(C) O2¯, S2¯, Cl(D) Mg2+, Ca2+, Sr2+
(E) Cs, Ba2+, La3+
CONSTRUCTED RESPONSE
22. Write both the long and short form of the ground state electron configuration for an arsenic
atom, showing the number of electrons in each subshell
1s2 2s22p6 3s23p63d10 4s24p3
[Ar] 3d10 4s24p3
23. The emission spectrum of hydrogen consists of several series of sharp emission lines in the
ultraviolet (Lyman series) in the visible (Balmer series) and in the infrared (Paschen series, Brackett
series, etc.) regions of the spectrum.
(a) What feature of the electronic energies of the hydrogen atom explains why the emission
spectrum consists of discrete wavelength rather than a continuum wavelength?
Quantized energy levels, Discrete energies, Wave properties of electron result in discrete
energy state.
(b) Account for the existence of several series of lines in the spectrum. What quantity
distinguishes one series of lines from another?
An electron in an excited-state atom can go to any of several lower energy states.
The lines in each series represent shifts from several higher energy states to a single
lower energy state, identified by the same principal quantum number or energy.
24. What characteristic of the electronic configurations of atoms distinguishes metals from
nonmetals? On the basis of this characteristic, explain why there are many more metals than
nonmetals.
Electron configurations: Metals: Valence electrons in s or d sublevels of their atoms. (A few
heavy elements have atoms with one or two electrons in p sublevels.) Nonmetals: Valence
electrons in the s and p sublevels of their atoms.
There are more metals than nonmetals because filling d orbitals in a given energy level
involves the atoms of ten elements and filling the f orbitals involves the atoms of 14
elements. In the same energy levels, the maximum number of elements with atoms receiving
p electrons is six.
25. Two important concepts that relate to the behavior of electrons in atom systems are the
Heisenberg uncertainty principle and the wave particle duality of matter.
(a) State the Heisenberg uncertainty principle as it related to the determining the position and
momentum of an object.
[either one of these]
(1) It is impossible to determine (or measure) both the position and the momentum of
any particle (or object or body) simultaneously.
(2) The more exactly the position of a particle is known, the less exactly the momentum
or velocity of the particle can be known.
(b) Explain why the uncertainty principle or the wave nature of particles is not significant when
describing the behavior of macroscopic objects, but it is very significant when describing
the behavior of electrons.
[either of these 2]
(1) The wavelength of a particle is given by the DeBroglie relation λ = h/mv. For masses of
macroscopic objects, h/m is so small for any v that λ is too small to be detectable. For an
electron, m is so small that h/mv yields a detectable λ.
(2) The product of the uncertainties in position and velocity depends on h/m and since h
is so small (h = 6.63x10-34 J/s), unless m is very small as with the electron, the product of
the uncertainties is too small to be detected.
26. Answer the following questions regarding light and its interactions with molecules, atoms,
and ions.
(a) The longest wavelength of light with enough energy to break the Cl–Cl boned in Cl2(g) is
495 nm.
(i) Calculate the frequency, in s–1, of the light.
v = c/ λ =
3.00 108 m/s
= = 6.06x1014 s-1
495 10–9 m
(ii) Calculate the energy, in J, of a photon of the light.
E = hv = (6.26x10–34 joule.s)(6.06x1014 s–1) = 4.02x10–19 J
(iii) Calculate the minimum energy, in kJ mol–1, of the Cl–Cl bond.
(4.02x10–19 J)(6.02x1023 mol–1) = 241,870 J = 242 kJ
(b) A certain line in the spectrum of atomic hydrogen is associated with the electronic
transition of the H atom from the sixth energy level (n = 6) to the second energy level (n =
2).
(i) Indicate whether the H atom emits energy or whether it absorbs energy during the
transition. Justify your answer.
emits energy. The n = 6 state is at a higher energy than n = 2, the electron must release
energy go to the lower state.
(ii) Calculate the wavelength, in nm, of the radiation associated with the spectral line.
1 1
E = –2.178x10–18 (22 – 62) J = 4.84x10–19 J
(iii) Account for the observation that the amount of energy associated with the same
electronic transition (n = 6 to n = 2) in the He+ ion is greater than that associated with
the corresponding transition in the H atom.
the He+ has a nuclear charge of 2+ vs H with a 1+, therefore, it has a stronger hold on the
electron. This requires more energy for the electron to move to higher energy levels and
when the electron moves from higher to lower energies, it releases more energy.
27. An AM radio station broadcasts at 1440 kHz and its FM partner broadcasts at 94.5 kHz
Calculate and compare the energy of photons emitted by these radio stations.
AM station E = hv = (6.26x10–34 joule.s)(1440 s-1) = 9.014 x10–31 J
FM station E = hv =(6.26x10–34 joule.s)(94.5 s-1) = 5.91 x10–32 J
Higher frequency = higher energy
28. Fireworks produce dramatic color effects by the use of metallic salts and highly exothermic
reactions. Barium salts are often responsible for the green color seen while calcium compounds
produce the orange color. The green color has a wavelength of 589nm while the orange is
approximately 628nm. The exothermic reaction causes some of the electrons in the respective
metal to become excited and move to higher energy levels. As these excited electrons lose
some or all of their extra energy, the energy is released in the form of visible light.
(A) Which color of light, green or orange, has a higher frequency?
Green has the higher frequency because it has a shorter wavelength
(B) Calculate the energy of a photon of green light.
E = h c/ λ
E = (6.626 x 10-34 Js)(2.998 x108 m/s) / (5.89 x10-7 m) == 3.37 x10-19 J
(C) Write the complete electron configurations for a neutral calcium atom and the shorthand
electron configuration for a neutral barium atom.
Ca = 1s22s22p6 3s23p64s3
Ba = [Xe] 6s2
(D) Considering their electron configurations, explain why calcium and barium produce
different flame colors.
Although they both have valence electron configurations that end with s 2