IB Chem I – Stations Review – Topic 2.2, 2.1
Name:_________________________________Partner
Name:_____________________ Date:__________
Work through the stations below with your partner in any order
you would like. Check your answers with the Station Keys. Then,
answer the Station Mini Quiz at the end. Plan to spend ~10 minutes
at each station. Work efficiently! Record your answers and create
your OWN question to leave behind for other groups! But don’t leave
them the answer. Extra points will be awarded to the best question
and you might see it on the test…
The review packet at the end of the stations is to be completed
outside of class for your own independent practice. Keys will be
posted on the website and it will not be collected. Please complete
the additional review packet by September 14th and bring any
questions you may have to go over during class before the review
game.
Keys can be found at http://shakeribchem.weebly.com
Go to Unit 1b – Topic 2 and scroll to the last links on the
page.
Unit 1b Exam Topic 2.1 / 2.2: September 15th, 2017 ONLY – no
extra time!
MC Section 20%
Short Answer & Problem Solving80%
ContentsPage #
Station 1 – Electron Configuration & The Rules2
Station 2 – Anomalies & Excited States4
Station 3 – Extension: Trouble with the Aufbau principle6
Station 4 – Electron Behavior, Energy, and Light7
Station 5 – Emission Spectra and Analysis9
Station 6 – Calculating Photon Energy11
Station 7 – The Bohr Model and Spectra12
Station 8 – Extension: Ionization Energy Predictions13
Station 9 – Average Atomic Mass and Isotopes14
Station 10 – Calculating Percentage Abundances16
Review Packet – Independent Practice17
(“Extensions” will not be tested on, but are helpful to deepen
our understanding and make excellent bonus questions.)Station 1 –
Electron Configuration & The Rules Write the full ground state
electron configurations and orbital notations (showing sublevels in
order of increasing energy) for the following:
# of e- Element (atom) e- configuration Orbital Notations/
diagrams
1._____lithium ________________________________
2._____calcium _______________________________
3._____nitrogen ______________________________
4._____chlorine ______________________________
Write the abbreviated ground state electron configurations for
the following:
# of electrons Element
5.______nitrogen ________________________________________
6.______chlorine ________________________________________
7.______iron ________________________________________
8. ______aluminum _______________________________________
Write a ground state electron configuration for these ions.
9.O2-:
______________________________________________________
10.K+:
________________________________________________________
11.Co3+:
______________________________________________________
For the following electron configurations determine the possible
elements (or possible ions) they may represent.
12.1s22s22p63s23p64s23d104p4
_________________________________
13. [Kr] 5s24d105p3
__________________________________________
Determine the errors in the configurations. Then, identify which
rule has been violated.
14.1s22s22p63s23p64s24d104p5 ____________________ Rule
violated:
15.1s22s22p63s33d5____________________ Rule violated:
16.[Ra] 7s25f8____________________ Rule violated:
17.[Kr] 5s24d105p5____________________ Rule violated:
Define the principles and rules.
18. Aufbau principle
19. Pauli-exclusion principle
20. Hund’s Rule
Check your answers and try the mini quiz at the station. Write
your question.
Mini quiz answers:
1. _______________________________________________________
2. _______________________________________________________Your
question & the correct answer:
Station 2 – Anomalies and Excited States Write the full ground
state electron configurations and orbital notations (showing
sublevels in order of increasing energy) for the following
anomalies:
# of e- Element (atom) e- configuration Orbital Notations/
diagrams
1._____copper ________________________________
2._____chromium _______________________________
Although the IB curriculum will not assess you on the following
exceptions, predict the ground state electron configurations for
the following metals that are also anomalies to the Aufbau
Principle. Write the abbreviated configuration and justify your
answer.
# of e- Element (atom) abbreviated e- configuration
3._____ Ag:
______________________________________________________
Why does this occur?
4._____ Mo:
______________________________________________________
Why does this occur?
Excited states of atoms can show electron behavior using
electron configurations.
5. How can an electron become “excited”?
Write an excited state full electron configuration for each.
6. Al
7. Ar
8. K
9. C
10. See the example, and determine which elements the excited
electron configurations represent.
a. 1s22s22p53s2 atom: ____Na___ what type of “jump” has
occurred? __n=2n=3_
b. 1s22s22p64s1 atom: __________ what type of “jump” has
occurred? ____________
c. 1s22s22p44s2 atom: __________ what type of “jump” has
occurred? ____________
Check your answers and try the mini quiz at the station. Write
your question.
Mini quiz answers:
1. _______________________________________________________
2. _______________________________________________________Your
question & the correct answer:
Station 3 – Extension: Trouble with the Aufbau principleRead the
article and consider the following questions.
1. Summarize the trouble with the Aufbau principle in 1-3
sentences.
2. Based on what we have studied about the anomalies and the
Aufbau principle and combining what you have read, why then do the
anomalies occur?
3. Why though, is it still useful to use the Aufbau
principle?
Check your answers and try the mini quiz at the station. Write
your question.
Mini quiz answers:
1_______________________________________________________
2_______________________________________________________Your
question & the correct answer:
Station 4 –Electron Behavior, Energy, and Light As you read
through the review text, answer the questions.
About 300 years ago, Sir Isaac Newton saw a beam of sunlight
through a glass prism. He discovered that light is made up of a
spectrum of seven distinct visible colors. This spectrum of colors
always appears in the same order. You can see this color spectrum
(Red, Orange, Yellow, Green, Blue, Indigo, Violet and all the
colors in between) when you look through a diffraction grating.
There are two color ranges that are not visible to our eyes in this
spectrum: below red is infra-red and above violet is
ultra-violet.
1. What is the wavelength range for:
a. Infrared ____________________ nm
b. Reds ____________________ nm
c. Oranges ____________________ nm
d. Yellows ____________________ nm
e. Greens____________________ nm
f. Blues____________________ nm
g. Violets ____________________ nm
h. UV____________________ nm
The color of a solid object depends on the colors of light that
it reflects. A red object looks red because it reflects red light
and absorbs all other colors. A blue object looks blue because it
reflects blue light and absorbs all other colors. A white object
reflects all colors of light equally and appears white. A black
object absorbs all colors and reflects no visible light and appears
black. Just like when you color with too many colors in one area
with crayons or markers, all colors are absorbed, none are
reflected and it appears black!
Type of Spectrum
Photographic example
Continuous (or continuum)
Absorption (dark line)
Emission (bright line)
2. So, emission spectra lines of color show
_____________________________ while absorption
spectra lines show____________________________.
Explanation of visible light at the electronic level:
What do fireworks, lasers, and neon signs have in common? In
each case, we see the brilliant colors because the atoms and
molecules are emitting energy in the form of visible light. The
chemistry of an element strongly depends on the arrangement of the
electrons. Electrons in an atom are normally found in the lowest
energy level called the ground state. However, they can be
"excited" to a higher energy level if given the right amount of
energy, usually in the form of 3. ______________________ or
_______________________. Once the electron is excited to a higher
energy level, it quickly loses the energy and "relaxes" back to a
more stable, lower energy level called the 4. ____________________
state. If the energy released is the same amount as the energy that
makes up visible light, the element produces a color. The visible
spectrum, showing the wavelengths corresponding to each color, is
shown below:
The important thing to know about absorption and emission lines
is that every atom of a particular element will have the same
pattern of lines all the time. And the spacing of the lines is the
same in both absorption and emission, only emission lines are added
to the continuum, while absorption lines are subtracted.
5. In the flame test, you observe a green light from a copper
compound, a lilac color from a potassium flame, and a deep red
flame from a lithium compound. Rank them in order of increasing
energy, and justify your answer.
4. Referring to question 5, which compound emitted the smallest
wavelength of light?
5. What is the relationship between energy, frequency, and
wavelength?
6. Do you think we can use the flame test to determine the
identity of unknowns in a mixture? Why or why not?
7. How are electrons “excited” in a flame test experiment?
8. What particles are responsible for the production of colored
light?
9. Why do different chemicals emit different colors of
light?
Check your answers and try the mini quiz at the station. Write
your question.
Mini quiz answers:
1_______________________________________________________
2_______________________________________________________Your
question & the correct answer:
Station 5 –Emission Spectra and Analysis Use the spectra at the
station to answer the following questions.
1. What do the different colors in a line spectrum represent?
Why are the spectra for each element unique?
2. It has been calculated that the observed colors in a hydrogen
atom correspond to the relaxation of the electron from a higher
energy level to the second energy level. Which color corresponds to
62 transition? 5 2 transition? 42 transition? 32 transition?
Explain your reasoning.
3. Which element produced the largest number of lines? Which
element produced the smallest number of lines? Explain why elements
produce different numbers of spectral lines.
4. What is the converging of spectral lines? Refer to an example
to explain.
Using the simplified spectra on the next page, answer the
questions.
5.List all elements present in unknown sample W.
6.List all elements present in unknown sample X.
7.List all elements present in unknown sample Y.
8.List all elements present in unknown sample Z.
Check your answers and try the mini quiz at the station. Write
your question.
Mini quiz answers:
1_______________________________________________________
2_______________________________________________________Your
question & the correct answer:
Station 6 –Calculating Photon Energy Max Planck theorized that
energy was transferred in
chunks known as quanta, equal to hv. The variable h is a
constant equal to 6.63 × 10 -34 J·s and the variable v (or
sometimes f) represents the frequency in s-1 or Hz (hertz). This
equation allows us to calculate the energy of photons, given their
frequency. If the wavelength is given, the energy can be determined
by first using the wave equation (c = λv) to find the frequency,
then using Planck’s equation to calculate energy.
Use the equations above to answer the following questions.
1. Ultraviolet radiation has a frequency of 6.8 × 1015 s-1.
Calculate the energy, in joules, of the
photon.
2. A sodium vapor lamp emits light photons with a wavelength of
5.89 × 10 -7 m. What is the energy
of these photons?
3. One of the electron transitions in a hydrogen atom produces
infrared light with a wavelength of
7.464 × 10 -6 m. What amount of energy causes this
transition?
4. What is the wavelength and frequency of photons with an
energy of 1.4 × 10 -21 J? Where does it fall on the electromagnetic
spectrum?
5. A photon emits a bright aqua light with a wavelength of 489
nm. What is the energy of this photon in Joules?
Check your answers and try the mini quiz at the station. Write
your question.
Mini quiz answers:
1_______________________________________________________
2_______________________________________________________Your
question & the correct answer:
Station 7 – The Bohr Model and Spectra Bohr applied Planck’s
concept of photon energy to the line spectra of elements. When
elements are excited they emit radiation at fixed wavelengths. He
proposed that only certain energy levels are allowed within the
structure of an atom. Electrons can move between these energy
levels. The light emitted by the elements is a measure of the
energy gap between the two electronic states.
1. From the image of the hydrogen electron behavior, what color
would we see for the electron jump n=1 n=2?
2.From the image of the hydrogen electron behavior, what color
would we see for the electron jump n=1 n=3?
3.From the image of the hydrogen electron behavior, what color
would we see for the electron jump n=1 n=4?
Extension: Relating Bohr’s model to photon energy. Use the
formulas, and try these!
6. Calculate the ΔE for the n = 3 to the n = 1 transition in
hydrogen. Where on the EMS would this appear? What does the sign
mean?
7. A hydrogen atom in its ground state absorbs light with a
wavelength of 102.6 nm. To which excited state would the electron
be excited? (Identify nf, the final energy level)
Check your answers and try the mini quiz at the station. Write
your question.
Mini quiz answers:
1_______________________________________________________
2_______________________________________________________Your
question & the correct answer:
Station 8 – Extension: Ionization Energy Predictions Ionization
energy is the energy required to completely remove an electron from
an atom. This can be thought of as the transition between n = 1 and
n = ∞. (∞ representing infinity, meaning the electron has been
completely removed)
1. How was our class definition of first ionization energy
different from the one provided?
2. Calculate the energy needed to remove the electron from
hydrogen in its ground
state.
3. What wavelength of light would be required to remove the
electron from
hydrogen in its ground state? Where is this on the EMS?
4.Write an ionization energy equation of hydrogen. (See your
notes.)
No mini quiz for this station.
Station 9 – Average Atomic Mass and Isotopes
1. The term “average atomic mass” is a
_________________________average, and so is calculated
differently from a “normal” average. Explain how this type of
average is calculated.
2. Define an isotope. Be specific.
3. How do the physical properties of isotopes differ? Be aware –
number of neutrons is NOT a testable physical property, but it does
influence the physical properties.
4. The element copper has naturally occurring isotopes with mass
numbers of 63 and 65.
The relative abundance and atomic masses are:
69.2% for mass of 62.93amu
30.8% for mass of 64.93amu.
Calculate the average atomic mass of copper.
5. Calculate the average atomic mass of sulfur if 95.00% of all
sulfur atoms have a mass of 31.972amu, 0.76%
has a mass of 32.971amu and 4.22% have a mass of 33.967 amu.
6. Naturally occurring strontium consists of four isotopes,
Sr-84, Sr-86, Sr-87 and Sr-88.
Below is the data concerning strontium:
Sr-84Sr-86Sr-87Sr-88
0.56%9.86%7.00% 82.58%
What is the average atomic mass of strontium?
7. Boron exists in two isotopes, boron-10 and boron-11. Based on
the atomic mass, which isotope should be more abundant?
8. Iodine is 80% 127I, 17% 126I, and 3% 128I. Calculate the
average atomic mass of iodine.
9. Hydrogen is 99% 1H, 0.8% 2H, and 0.2% 3H. Calculate its
average atomic mass.
Check your answers and try the mini quiz at the station. Write
your question.
Mini quiz answers:
1_______________________________________________________
2_______________________________________________________Your
question & the correct answer:
Station 10 – Calculating Percentage Abundances (Remember,
percentages must be in decimal form during your calculations…)
1. Bromine has two naturally occurring isotopes. Bromine-79 has
a mass of 78.918 amu and is 50.69% abundant. Using the atomic mass
reported on the periodic table, determine the mass of bromine-81,
the other isotope of bromine. Show all work.
2. Antimony has two naturally occurring isotopes. The mass of
antimony-121 is 120.904 amu and the mass of antimony-123 is 122.904
amu. Using the average mass from the periodic table, calculate the
abundance of each isotope. Show all work.
4.Strontium consists of four isotopes with masses of 84
(abundance 0.50%), 86 (abundance of 9.9%), 87 (abundance of
unknown), and 88 (abundance of unknown). Using the average atomic
mass of strontium, calculate the percentage abundances of the
heaviest two isotopes. Show all work.
Check your answers and try the mini quiz at the station. Write
your question.
Mini quiz answers:
1_______________________________________________________
2_______________________________________________________Your
question & the correct answer:
Review Packet – Independent Practice
Check your answers at http://shakeribchem.weebly.com
#eAtome configurationorbital diagram
1._____ oxygen ________________________________
_________________________________
2._____potassium ____________________________
_________________________________
3._____hydrogen _____________________________
__________________________________
4. _____ neon __________________________________
________________________________
5. _____ phosphorous ___________________________
_________________________________
Abbreviated configurations
6.______zinc ________________________________________
7.______barium ________________________________________
8. ______bromine ________________________________________
9. ______magnesium _______________________________________
10. ______fluorine
__________________________________________
Electron Configuration Elements (atoms) and Ions
Write the electron configuration and orbital notations for the
following Atoms and ions:
Element / Ions
Atomic number
# of e-
Electron Configuration/Orbital Diagrams
F1-
Na1+
Al3+
Cl1-
Br1-
Mg2+
E.In the space below, write the full (unabbreviated) electron
configurations of the following elements:
1)sodium________________________________________________
2)iron ________________________________________________
3)bromine ________________________________________________
4)barium ________________________________________________
5)neptunium ________________________________________________
F.In the space below, write the Noble Gas (abbreviated) electron
configurations of the following elements:
6)cobalt________________________________________________
7)silver________________________________________________
8)tellurium________________________________________________
9)radium________________________________________________
10)lawrencium________________________________________________
11)manganese________________________________________________
12)silver________________________________________________
13)nitrogen________________________________________________
14)sulfur________________________________________________
15)argon________________________________________________
G. Given an element with atomic number 11, provide the following
information:
a. How many electrons will fill each of the following shells and
list their subshells.
1st shell:
2nd shell:
3rd shell:
b. Is this element likely to form a cation or
anion?_________
c. How do you know?
d. What charge will the ion formed by this element have?
H. Explain, based on electron configuration, why the noble gases
are so unreactive. Use helium and neon as examples to illustrate
your explanation.
I. Determine which of the following electron configurations are
not valid, and why.
a. 1s22s22p63s23p64s24d104p5
b. 1s22s22p63s33d5
c. [Ra] 7s25f8
d. [Kr] 5s24d105p5
e. [Xe]
J. Is light a particle or a wave?
Is light composed of waves or of particles? If light is waves,
then one can always reduce the amount of light by making the waves
weaker, while if light is particles, there is a minimum amount of
light you can have - a single ``particle'' of light. In 1905,
Einstein found the answer: Light is both! In some situations, it
behaves like waves, while in others it behaves like particles.
This may seem odd. How can light act like both a wave and a
particle at the same time? Consider a duck-billed platypus. It has
some duck-like properties and some beaver-like properties, but it
is neither. Similarly, light has some wavelike properties and some
particle like properties, but it is neither a pure wave nor a pure
particle.
A wave of light has a wavelength, defined as the distance from
one crest of the wave to the next, and written using the symbol .
The wavelengths of visible light are quite small: between
400 mm and 650 nm, where
1 nm = 10-9 m is a ``nanometer'' - one
billionth of a meter. Red light has long wavelengths, while blue
light has short wavelengths.
A particle of light, known as a photon, has an energy E. The
energy of a single photon of visible light is tiny, barely enough
to disturb one atom; we use units of “electron-volts”, abbreviated
as eV, to measure the energy of photons. Photons of red light have
low energies, while photons of blue light have high energies.
The energy E of a photon is proportional to the wave frequency
v
E = h v
where the constant of proportionality h is the Planck's
Constant, h = 6.626 x 10-34 J s.
Also, the relationship between frequency and wavelength can be
defined as:
v= c
λ
where c is the speed of light (3×108 metres per second).
So, photons still have a wavelength. A famous result of
quantum mechanics is that the wavelength relates to the energy of
the photon. The longer the wavelength, the smaller the
energy. For instance, ultraviolet photons have shorter
wavelengths than visible photons, and thus more energy. This
is why they can give you sunburn, while ordinary light cannot.
One means by which a continuous spectrum can be produced is by
thermal emission from a black body. This is particularly relevant
in astronomy. Astronomical spectra can be combination of absorption
and emission lines on a continuous background spectrum.
One convenient method of exciting atoms of an element is to pass
an electric current through a sample of the element in the vapor
phase. This is the principle behind the spectrum tubes. A spectrum
tube contains a small sample of an element in the vapor phase. An
electric discharge through the tube will cause the vapor to glow
brightly. The glow is produced when excited electrons emit visible
light energy as they return to their original levels.
When visible light energy from a spectrum tube is passed through
a diffraction grating, a bright line spectrum, or line-emission
spectrum is produced. Each element has its own unique emission
spectrum by which it can be identified, analogous to a fingerprint.
Such a spectrum consists of a series of bright lines of definite
wavelength. Each wavelength can be mathematically related to a
definite quantity of energy produced by the movement of an electron
from one discrete energy level to another. Thus, emission spectra
are experimental proof that electrons exist in definite,
distinctive energy levels in an atom.
Refer to this text to answer the following questions.
1. What is the difference between a line spectrum and a
continuous spectrum?
2. Each line in the emission spectrum of the hydrogen
corresponds to an electromagnetic radiation with a specific
wavelength. Match the 4 observed colors with the following
wavelengths: 410 nm, 434 nm, 486 nm, and 656 nm.
3. How are electrons “excited”? What happens when the electrons
“relax”?
4. Each element has its own unique line emission spectrum, just
like fingerprints. Explain how this technique can be used to
determine the elemental composition of stars.
5. How can the difference in the brightness of spectral lines be
explained?
6. According to the modern theory of the atom, where may an
atom’s electrons be found?
7. State the equation used to determine the energy content of a
packet of light of specific frequency.
8. What form of energy emission accompanies the return of
excited electrons to the ground state?
9.Explain, in terms of electron transition, how bright-line
spectra are produced by atoms.
K. Solving for photon energy
1. Find the energy, in joules per photon, of microwave radiation
with a frequency of 7.91 × 10 10 s-1.
2. Find the energy in kJ for an x-ray photon with a frequency of
2.4 × 10 18 s-1.
3. A ruby laser produces red light that has a wavelength of 500
nm. Calculate its energy in joules.
4. What is the frequency of UV light that has an energy of 2.39
× 10 -18 J?
5. The frequency of violet light is about 7.495x10 14 Hz. What
is the wavelength of this radiation?
6. What is the frequency of a photon that has a wavelength of
1428 nm? What type of radiation is this?
7. A popular radio station broadcasts at 107.9 MHz (M = 10 6 ).
Find the wavelength of this radiation, in meters, and the energy of
one of these photons, in J. What type of radiation is this?
8. What is the energy of a photon with:
a) a wavelength of 58 nm? What type of radiation is it?
b) a wavelength of 0.065 cm? What type of radiation is it?
9. Which of the following are directly related?
a) energy and wavelength
b) wavelength and frequency
c) frequency and energy
L. Planck recognized that energy is quantized and related the
energy of radiation
(emitted or absorbed) to its frequency.
ΔE = n h ν where n = integer and h = Planck's constant = 6.626 x
10 -34 J s
1. What is the energy needed to remove the remaining electron
from He+ in its ground state? Is it easier or harder to remove the
electron from He+ than from H?
M. Average Atomic Mass
1. Calculate the average atomic mass of bromine. One isotope of
bromine has an atomic mass of 78.92amu and
a relative abundance of 50.69%. The other major isotope of
bromine has an atomic mass of 80.92amu and a
relative abundance of 49.31%.
2. Lithium-6 is 4% abundant and lithium-7 is 96% abundant. What
is the average mass of lithium? Try this WITHOUT a calculator!
3.Here are three isotopes of an element: 12C13C 14C
a.The element is: __________________
b.The number 6 refers to the _________________________
c.The numbers 12, 13, and 14 refer to the
________________________
d.How many protons and neutrons are in the first isotope?
_________________
e.How many protons and neutrons are in the second isotope?
_________________
f.How many protons and neutrons are in the third isotope?
_________________
4.Complete the following chart:
Isotope name
atomic #
mass #
# of protons
# of neutrons
# of electrons
potassium-37
oxygen-17
uranium-235
uranium-238
boron-10
boron-11
5. Argon has three naturally occurring isotopes: argon-36,
argon-38, and argon-40. Based on argon’s reported atomic mass,
which isotope exist as the most abundant in nature? Explain without
calculations.
6. Calculate the atomic mass of lead. The four lead isotopes
have atomic masses and relative abundances of 203.973 amu (1.4%),
205.974 amu (24.1%), 206.976 amu (22.1%) and 207.977 amu
(52.4%).
7. Titanium has five common isotopes: 46Ti (8.0%), 47Ti (7.8%),
48Ti (73.4%),
49Ti (5.5%), 50Ti (5.3%). What is the average atomic mass of
titanium?
8. What is the atomic mass of hafnium if, out of every 100
atoms, 5 have a mass of 176, 19 have a mass of 177, 27 have a mass
of 178, 14 have a mass of 179, and 35 have a mass of 180.0?
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