Download Chemistry I Accelerated StudyGuideline

Chemistry I Accelerated
Study Guideline - Chapter Five
Electrons in Atoms
____________________________________________________
By the end of this chapter, the skills you should be able to demonstrate
are:
1. Use the Bohr model to explain atomic emission spectrum.
2. Explain the relative wavelength, frequency and energy of the eight
parts of the electromagnetic spectrum.
3. Calculate wavelength, frequency or energy when given one of the
three.
4. Describe the contributions of Bohr, Heisenburg, de Broglie, and
Schrödenger to quantum theory.
5. Descibe wave-mechanical model of the the atom.
6. Explain the theory of atomic emission spectra using quantum theory.
7. Describe an electron cloud.
8. Compare Dalton’s, Thomson’s, Bohr’s and the
de Broglie/Schrödinger’s models of the atom.
9. Explain the significance of quantized energies and the quantum
mechanical model of the atom.
10. Characterize the four quantum numbers.
11. Use the Pauli exclusion principle and quantum numbers to describe
an electron in an atom.
12. Describe the general shape of s, p and d orbitals.
13. Apply the Aufbau principle, Pauli exclusion principle, and Hund’s
rule in writing electron configurations for atoms.
14. Explain why the electron configurations for Cr and Cu differ from
those assigned using the standard protocol.
15. Write electron dot diagrams for elements.
Suggested Problems: p.149-151: #26, 40, 48, 54, 56, 74
page 1
Bill Nye - Waves
PLEASE ANSWER THE FOLLOWING QUESTION WHILE YOU WATCH THE VIDEO
1. Define the following terms:
Wavelength
Frequency
Amplitude
2. How does energy travel?
3. What is the relationship between frequency and wavelength?
4. What is different about each color of the spectrum?
5. What are some examples of electromagnetic waves?
6. What is the relationship between frequency and energy?
page 2
Waves
Light travels through space by means of waves. Each wave has a frequency (n), a wavelength (l) and an
amplitude. The figure below represents a light wave. Label the wavelength and the amplitude.
1. What is the wavelength (in meters) of the wave?
2. If the frequency is 6.00 x 101 4/sec what is the product of the wavelength times the frequency?
3. What is the significance of the value just calculated?
4. Write a mathematical expression of what you have just found.
5. The light in the figure corresponds to green light. Describe how the value of the wavelength would differ if
the light were red.
6. How would the frequency of red light compare to that of green light?
7. How would the value of the product of the wavelength times the frequency for red light compare to that of
green light? Explain.
8. How do the energies of green and red light compare? What equation expresses this relationship?
page 3
Electromagnetic Radiation
1. The speed of light in a vacuum is _______________ meters per second.
2. All waves can be described in terms of their amplitude, frequency and ____________________.
3. Early in the last century, scientists found that light has the characteristics of both waves and
____________.
4. The ______________ of a wave is the number of complete waves passing a fixed point in a given time.
5. The wavelength of microwave radiation is ______________ than the wavelength of visible light.
6. The color of visible light that has the longest wavelength is ______________.
7. A heat lamp produces ______________ radiation.
8. A wave with a high frequency has a ______________ wavelength.
9. What is the frequency of a beam of blue light with a wavelength of 595 nanometers?
Type of Radiation
Description of Wave
1. _________________________________
2. _________________________________
3. _________________________________
These waves have a long
wavelength and a low
frequency
4. _________________________________
5. _________________________________
These are the colors of the
visible spectrum (wavelengths
between 750 and 400 nm).
6. _________________________________
7. _________________________________
Increasing
Wavelength 8. _________________________________
page 4
These waves have a short
wavelength and a high
frequency.
Wavelength/Frequency/Energy Practice Problems
h = 6.626 X 10–34 J•s
c = 3.00 X 108 m/s
E = hν
14
1. What is the energy of light with a frequency of 4.31 X 10
λ =c/ν
Hz?
2. A certain violet light has a wavelength of 413 nm. What is the frequency and the energy of the light?
–14
3. Gamma rays have a wavelength of 5.00 X 10
rays?
meters. What is the energy and the frequency of gamma
4. Public Radio in NYC broadcasts at 820 Kilohertz. What is the energy and wavelength of their broadcast?
page 5
Light and Energy Levels of the Atom
The quantum level occupied by an electron in an atom depends on the energy of the electron. Changes in
quantum level are related to the absorption or emission of energy. The figure below represents the four
lowest energy levels of an atom (n = 1 to 4). The six lettered arrow represent changes in the energy level of
an electron.
4
3
A
B
C
D
E
F
2
1
1. Which of the lettered energy changes involve the absorption of energy by the atom?
_____________
2. Which of the lettered energy changes involve the emission of light energy by the atom?
_____________
3. Of the three lettered energy changes that involve emission one results in emission of blue light, one results
in emission of yellow light, and one results in emission of UV light. Which one is blue __________?
Which one is yellow __________? Which one is UV __________?
Quantum Theory
If the statement is true write true. If it is false, change the underlined word(s) to make the staement true.
___________ 1. We are not aware of quantum effects in the world around us because quanta of light
energy are very large.
___________ 2. In sodium metal, violet light causes the photoelectric effect but red light does not because
photons of violet light have less energy than those of red light.
___________ 3. In the photoelectric effect, protons are ejected from the surface of a metal when light
shines upon it.
___________ 4. We are constantly surrounded by low frequency X-rays.
___________ 5. Plank’s theory relate frequency of radiation to its energy.
___________ 6. The wavelengths of radiation emitted by a hot object shift as its temperature increases.
___________ 7. The dual nature of light means that light has the properties of a charge and a wave.
8. Explain using a set if stairs and a ramp the difference between a continuous change and a quantized
change.
page 6
The Modern Atom
1. Every element has a uniquely characteristic atomic _____________ spectrum.
2. Heisenberg’s uncertainty principle states that the position and _______________ (speed in a direction) of
an electron can not be measured and known at the same time.
3. ___________ used Balmer’s idea of energy quantization to explain the line spectrum of hydrogen.
4. Each __________ _________ is labeled with a primary quantum number.
5. An electron that absorbs a quantum of energy jumps to a level of _______________ energy called an
excited state.
6. _______________ is emitted when an electron falls back to a lower energy level.
7. A(n) _______________ is a region in space where an electron with a particular energy is likely to be
found.
8. Each electron is described by _______________ quantum numbers.
9. The ________ ___________ _________ states that each orbital can only hold two electrons and that
these electrons must have opposite spins.
10. The modern model of the atom is called the ____________–____________ model.
11. The Bohr model is inaccurate because each electron has no definite _______________.
12. De Broglie derived a mathematical relationship between mass and velocity of a moving particle and
showed that the particle would also show ________–________ behavior.
13. Sodium emits _______________ light.
14. An electron cloud is an area where there is the _______________ likelihood of finding an electron.
15. The first orbital in each principle energy level is a(n) _______________.
16. P orbitals are _______________ shaped and have _______________ lobes.
17. The 3s orbital differs from the 2s orbital in that it is _______________.
18. The number of sublevels in each principle energy level equals the number of the ______________
____________.
19. The f sublevel is first found on the _______________ energy level.
Sublevel
Number of Orbitals
Maximum Number of Electrons
s
p
d
f
page 7
Electron Configurations
1. Identify the element that has each of the following electron configurations:
2
2
3
2
2
6
2
3
2
2
6
2
6
2
1
2
2
6
2
6
2
10
2
2
6
2
6
2
10
a. 1s 2s 2p
________
b. 1s 2s 2p 3s 3p
________
c. 1s 2s 2p 3s 3p 4s 3d
d. 1s 2s 2p 3s 3p 4s 3d
e. 1s 2s 2p 3s 3p 4s 3d
________
6
2
10
6
2
10
4p 5s 4d
4p 5s 4d
________
6
2
5p 6s 4f
14
6
5d
________
2. Write out the electron configuration for each of the elements. Note how many paired and unpaired
electrons each has.
a. lithium
electron configuration __________________________________________________________
number of paired electrons
________ number of unpaired electrons
________
b. potassium
electron configuration __________________________________________________________
number of paired electrons
________ number of unpaired electrons
________
c. silicon
electron configuration __________________________________________________________
number of paired electrons
________ number of unpaired electrons
________
d. cobalt
electron configuration __________________________________________________________
number of paired electrons
________ number of unpaired electrons
________
e. sulfur
electron configuration __________________________________________________________
number of paired electrons
________ number of unpaired electrons
________
f. vanadium
electron configuration __________________________________________________________
number of paired electrons
________ number of unpaired electrons
________
g. selenium
electron configuration __________________________________________________________
number of paired electrons
________ number of unpaired electrons
page 8
________
h. mercury
electron configuration __________________________________________________________
__________________________________________________________
number of paired electrons
________ number of unpaired electrons
________
i. bismuth
electron configuration __________________________________________________________
__________________________________________________________
number of paired electrons
________ number of unpaired electrons
________
j. mendeleevium
electron configuration __________________________________________________________
__________________________________________________________
number of paired electrons
________ number of unpaired electrons
3. Draw the orbital diagram for each of the following:
a. Oxygen
b. Neon
c. Magnesium
d. Potassium
e. Aluminum
f. Phosphorus
g. Argon
page 9
________
h. scandium
i. Iodine
3+
j. B
2+
k. Si
l. As
3–
–
m. Br
True/False: If the statement is true write “True”; if it is false change the underlined word(s)
to make the statement true.
_________ 1. The Pauli exclusion principle states that an orbital can hold a maximum of two electrons
_________ 2. The sum of the superscripts in an electron configuration represents the total number of
neutrons in the atom.
_________ 3. The Aufbau principle states that electrons are added one at a time to the highest energy
orbitals available until all the electrons of the atom have been accounted for.
_________ 4. An orbital diagram uses arrows to represent the spin of the electrons.
_________ 5. The ground state is the least stable energy state of the atom.
_________ 6. According to Hund’s rule, electrons occupy equal energy orbitals so that a maximum
number of unpaired electrons results.
page 10
Challenge Problems - Electrons
1. In a photoelectric effect experiment a student shines a light of greater than the threshold frequency upon
the surface of the metal. She observes that after a long time the number of ejected electrons begins to
decline. Can you explain why?
2. Look up the electron configuration of rhodium. Why is it exceptional? Explain.
3. Draw the orbital diagram for copper.
4. How many lobes will the f orbitals have? Explain.
page 11
Atoms and Waves Crossword
ACROSS
1. Mathematical or Chemical expression
5. The last element in the righthand column of
the periodic table
6. Element that has 16 protons and 16 electrons
9. Energy needed to remove an electron from an
atom
16. Height of a crest
17. Symbol for an element that has 22 protons
and 22 electrons
18. Number of electrons in an H+ ion
19. As wavelength increases, frequency and energy
______________
20. For an atom to become positively charged, it
must ___________ electrons
21. Without net charge
22. Number of electrons in a neon atom
24. Negatively charged subatomic particle
25. Is made up of photons
DOWN
2. Discreet packets of or levels of energy
3. Total number of electrons that can be held in
the 3d sublevel of an atom
4. Greek letter symbol for frequency
6. The nearest star
7. Number of waves per unit time
8. Color of light which has the highest energy
10. Wave location for which amplitude is zero
11. Blue stars have a higher value of this than do
red stars
1
12. Region of space in which an electron is
likely to be found
13. An atom in the lowest energy state is in the
___________ state
14. Distance between crests
15. Device that produces a beam of electrons:
cathode-ray-___________
20. An energy state is also called an energy ______
22. Number of electrons that can exist in an orbital
23. Same as 4 Down
2
3
4
5
6
7
8
9
10
11
12
14
13
15
16
17
18
19
20
21
22
24
25
page 12
23