Download Atomic Structure PSI CW and HW Packet answer key edited STIRES

Atomic Structure: PSI CW & HW Packet
Bohr Model
Classwork
1. Describe Rutherford’s nuclear model of the atom.
In the nuclear model of the atom, protons (and neutrons) are housed in a small, dense
nucleus. Electrons surround the nucleus in an area of mostly empty space.
2. Explain the problems with Rutherford’s nuclear model of the atom.
Rutherford’s model could not explain why most atoms are stable and why excited atoms
produce an emission spectrum.
3. According to Niels Bohr, what does “n” stand for?
According to Bohr, “n” stands for the orbit or energy level of the electron.
4. According to Bohr, why do atoms emit light?
According to the Bohr model, atoms emit light because excited electrons are returning to
lower energy states, emitting the energy difference. This energy always has a specific
wavelength because the electrons can only exist in set orbits.
5. How do electrons get to an excited state?
Electrons absorb energy is quantized amounts.
6. Which of the following transitions would produce the greatest amount of energy: 21, 32, or
43?
21
7. Define the following terms: emission spectrum and absorption spectrum
Emission – frequencies of light released Absorptions – frequencies of light absorbed
8. How can emission spectra be used to identify an unknown element?
Because each element has a different number of protons and electrons, they’re “n”
values differ from each other. Each element thus produces its own emission spectra
which allows scientists to identify elements.
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Homework
9. How did emission spectra of gases affect our understanding of atomic structure?
Emission spectra showed that electrons only emit radiation at certain wavelengths and
frequencies, and, therefore, energy levels. This indicating that electrons could be found in
specific energy levels or orbits.
10. Describe the Bohr model of an atom.
In the Bohr model of the atom, the nucleus is at the center of the atom with electrons
orbiting it at set distances, similar to the way planets orbit the Sun. The different orbits
correspond to different energy levels. Rutherford’s nuclear model did not address the
energy of electrons.
11. What element was the Bohr model based on?
Hydrogen (1 protons and 1 electron)
12. Define “ground state”.
The ground state is the lowest energy state that an electron can reside.
13. Define “excited state”.
Excited states are any energy state above the ground state.
14. Which of the following transitions would produce the greatest amount of energy: 12, 31, or
21?
3 1
15. Describe why different electron transitions produce different colors.
The energy differences between levels will give off different amounts of energy. The
different frequencies associated with the energy given off will produce certain colors.
16. What do lines on an emission spectrum correspond to, according to Bohr?
Electrons moving from higher energy levels to lower levels release energy. The lines
correspond to the wavelengths of light energy released.
17. Why did scientists need to improve upon the Bohr model?
The Bohr model was only able to explain the emission spectrum of hydrogen.
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Quantum Mechanics
Classwork
18. What is the Heisenberg Uncertainty Principle?
The position and momentum of an electron cannot be known simultaneously.
19. What do quantum numbers tell us about electrons in an atom?
The probability of finding an electron in a specific orbital.
Homework
20. How does the quantum mechanical model treat an electron?
The quantum mechanical model treats the electron as both a wave and a particle. One can
only approximate the probable location of finding an electron.
21. Why does the dual nature of matter make it difficult to observe very small particles like
electrons?
Observing these particles changes their momentum, so position and momentum cannot be
simultaneously known.
The Quantum Model
Classwork
22. Describe the quantum mechanical model.
Electrons are treated as both a wave and a particle in the quantum mechanical model.
The location of an electron can only be approximated, according to the quantum
mechanical model.
23. Name and describe the 4 quantum numbers.
Principal quantum number (n) – energy level of the electron; angular quantum number (l) –
shape of the orbital; magnetic quantum number (ml) – orientation of the orbital; spin
quantum number (ms) – direction of the electron spin
24. Give the number of orientations for each type of orbital.
S orbitals – 1 orientation; p orbitals – 3 orientations; d orbitals – 5 orientations; f orbitals –
7 orientations
25. What orbital shapes can be found in the n = 1 level?
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Only s
26. What is the maximum number of electrons that could be in the n = 3 level?
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27. The spin quantum number has how many possible values?
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Homework
28. Give the maximum number of electrons in s, p, d, and f orbitals.
S – 2; p – 6; d – 10; f – 14
29. What subshells can be found in the n = 3 level?
3s, 3p, 3d
30. What is the maximum number of electrons that could be in the n = 2 level?
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31. Which subshell has the higher energy: 3d or 4s?
3d
Electron Configurations
Classwork
32. Draw the energy level diagram for Iron.
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33. Draw the energy level diagram for Sulfur.
34. Draw the energy level diagram for Argon
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35. Draw the energy level diagram for Neon.
36. What is the electron configuration of Iron?
1s2 2s2 2p6 3s2 3p6 4s2 3d6
37. What is the electron configuration of Bromine? 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p5
38. What is the electron configuration of Lithium? 1s2 2s1
39. Identify the following element: 1s22s22p63s23p2 Silicon
40. Can an electron configuration be 1s32s1? Explain.
No, it violates the Pauli Exclusion Principle
Homework
41. Draw the energy level diagram for Titanium.
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42. Draw the energy level diagram for zinc.
43. Draw the energy level diagram for Krypton.
44. What is the electron configuration of Strontium? 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2
45. What is the electron configuration of Nickel? 1s2 2s2 2p6 3s2 3p6 4s2 3d8
46. What is the electron configuration of Francium?
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10 6p6 7s1
47. Identify the following element: 1s22s22p63s23p24s23d104p65s1 Rubidium
48. Why is it incorrect to write the electron configuration of selenium, 1s22s22p63s23d103p24s24p4
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Electron configurations are written in terms of increasing energy levels, 3d has a higher
energy level than 4s
Free Response
1. The currently accepted best model of the atom is the quantum mechanical model. Trace the
evolution of this atomic model by describing each of the following and the problems that were
resolved in the next model.
a. Plum Pudding Model
The plum pudding model of the atom consisted of a positively charge sphere with
electrons embedded in it. It was developed because it was support by the finding
that protons were much more massive than electrons and by Coulomb’s law of
attraction between oppositely charged particles. It was disproved by the gold foil
experiment which demonstrated that the atom was mostly empty space with a
small dense core.
b. Nuclear Model
The nuclear model consisted of an atom of mostly empty space with a dense core
containing the protons and neutrons. It could not explain the lack of continuous
spectrum of energy being emitted from the atom.
c. Bohr Model
The Bohr model explained the quantized spectra lines as set orbits where the
electrons could exist without emitting energy, but it could not explain why the
electrons did not emit energy in these orbits.
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d. Quantum Model
The quantum model is based on the wave-particle duality of matter and can only
be used to determine the probable path and location of electrons but not their
exact orbit or location. It cannot explain why or how an electron can behave as
both a particle and a wave.
2. The emission spectra of three elements is given below. Use these spectra to answer the
following questions.
a. Explain why these emission spectra occur.
According to the Bohr model, emission spectra occur in quantized amounts because
electrons orbit the nucleus at set distances and when they fall from higher to lower
energy orbits they emit light at a set wavelength.
b. What is the significance of the hydrogen spectrum?
The hydrogen spectrum was the first and only spectrum that could be adequately
explained by the Bohr model.
c. Draw the energy level diagram for hydrogen and sodium.
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3. The emission spectra of two elements is given below. Use these spectra to answer the
following questions.
a. How would the Bohr model explain the greater number of spectral lines in neon vs.
helium?
Neon has more electrons it is expected to have more energy levels/orbits.
b. Write the electron configuration for these elements.
Helium: 1s2 ; Neon: 1s22s22p6
4. Carbon-14 is an unstable isotope of carbon with a half-life of 5730 years.
a. Draw the energy level diagram for this element. Explain how you used Hund’s Rule, the
Aufbau principle, and the Pauli Exclusion Principle to construct your diagram.
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Pauli Exclusion principle only allows
you to put two electrons in each
orbital, drawing one arrow up and
the other down. The Aufbau
principle says to fill the 1s before
the 2s, and then move to 2p. Hund’s
Rule says to put one arrow
(electron) in each 2p space before
doubling up
b. Write the electron configuration for carbon-14.
1s22s22p2
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