Download Unit 10 Student Packet PS and FRQs

AP Chemistry: Unit 10 Atomic Structure Reading Assignment
Part 1: Atomic Structure Summary – please also reference the textbook for information!
Our model of the atom continues to evolve as new discoveries are made. The first atomic model
that was based on scientific experiments came from John Dalton. He believed that each element had a
smallest subunit which he called the atom. He believed the atom could not be subdivided into smaller
parts. Later, we learned, his model was inadequate. After J.J Thompson’s experiments with cathode ray
tubes, the electron was discovered and after Ernest Rutherford’s gold foil experiment it became clear
that an atom consisted of a very tiny, dense, positively charged nucleus surrounded by an electron cloud.
Later, after searching in vain for a positively charged particle nearly the same size as an electron,
Rutherford’s experiments led him to the fact that the positively charged particle he was searching for,
eventually named the proton, is 1840 times more massive than the electron even though it carries the
same amount of charge.
Before the nucleus was discovered, scientists had noticed and measured radioactivity in atoms.
They observed that a sample of a pure substance like Radon would give off large amounts of energy and
then, somehow, become impure (mixed with other elements). They had also measured the relative
masses of the atoms but could not figure out what made up an atom’s mass since they knew the number
of protons and electrons must be equal, but electrons were too tiny to really contribute to the mass. In
addition, they noticed that some atoms had the same properties as others but different masses. Once
the neutron was discovered in 1932 and once scientists put together Einstein’s idea of E=mc2, these
observations could be explained.
Some atoms have isotopes which means they have the same number of electrons and protons but
different numbers of neutrons. This gives them different masses (the mass of the atom comes mostly
from protons and neutrons since electrons are 1/1840 the size of a proton), but the same chemical and
physical properties because they have the same number of electrons and protons. Radioactivity is
caused by unstable nuclei that are too large or have too many or too few neutrons. Unstable nuclei split
apart into smaller atoms which is why a pure sample becomes impure after sitting for awhile. When the
nuclei split, some of the mass is converted to energy – an amount of energy that can be calculated from
Einstein’s equation E=mc2, and this is where the radiation observed was coming from.
Because a given element can have multiple isotopes, it becomes important to be able to distinguish
between them. We do this using isotope notation by either writing the “symbol-mass number” as in C-12
or C-14, or by writing the mass number as an exponent to the left of the atom’s symbol as in 14C or 146C,
where the atomic number of the atom is the subscript to the left of the symbol.
Also, because isotopes exist, it is important to know that the masses listed on the periodic table
are weighted averages. For example, Carbon has 2 isotopes, C-12 and C-13 with masses of 12 amu and 13
amu respectively. (Remember that amu stands for atomic mass unit and is a unit that scientists made up
since they couldn’t weigh individual atoms. Since then we’ve been able to determine that 1 amu = 1.66
x10-27 Kg). The most abundant carbon isotope is Carbon-12 since it makes up 98.93% of the carbon on
earth. The other 1.07% is the carbon-13. In order to use these masses in a practical way (solving
stoichiometry problems), we need to have an accurate way of knowing the mass of 1 mole of each
element. We must take into consideration the abundance of a particular isotope when calculating the
average atomic mass. Since only 1.07%, in carbon’s case, will contribute a heavier mass to the average.
To calculation average mass use the following equation:
(mass1)(abundance1) + mass2(abundance2)+… = average atomic mass
(Don’t forget to divide the percentages by 100).
Review of Chapter 2 Questions
1. Describe the experimental evidence that lead to the discovery of the electron and the
charge:mass ratio for an electron.
2. Describe Thomson’s model of the atom.
3. Describe Bohr’s atomic model and how it relates to emissions spectra (may need to look in another
chapter).
4. Describe the experimental evidence that enabled Robert Millikan to determine the magnitude of
the electron charge and to calculate the mass of an electron.
5. Describe the experimental evidence that lead to the discovery of the nucleus.
6. Describe the nuclear atom.
7. Calculate the average atomic mass given the following isotopes and their percent abundances.
a) Boron-10 19.78%; Boron-11 80.22%
b) Te-122 0.089%; Te-124 2.46%; Te-126 5.48%; Te-128 91.97%
8. There are three isotopes of element X. X-45, X-46, and X-48. If the atomic mass of element X is
45.23 which isotope is most common? Why?
9. What is the difference between Copper-65 and Copper-63? Do they have the same chemical and
physical properties?
10. Pg 82 # 101 & 104
Part 2: Intro to Chapter 7 Reading
*** Read the study guide that I have posted online for help and a great summary of Ch 7!
Niels Bohr used atomic emission spectra to change the model of the atom from one with a nucleus
and undefined electron cloud, to an atom with a nucleus and distinguishable energy levels.
The trends can be explained through atomic structure. Recall that an atom has a tiny nucleus in the
center where the positive charge of the atom is located. Electrons exist in the space around the nucleus
in energy levels that make up the electron cloud. Scientists don’t really know what holds the electrons in
a particular energy level, nor what prevents them from falling into the nucleus. However, it is useful to
think of the electrons as attracted to the nucleus and repelled by electrons close by. As atomic
number increases, the positive charge in the nucleus increases and electrons, within the same energy
level, are more strongly attracted to the nucleus. The electrons are, therefore, closer to the nucleus
and more strongly held (For you physics students: remember that the electrostatic force, F=kq1q2/r2
increases as r, radius, decreases). As energy level increases, valence electrons are less tightly held to
the nucleus for two reasons, they are farther from the nucleus and some of the attractive force of the
nucleus is cancelled out by the inner electrons. We say that the valence electrons are shielded by inner
electrons. The more energy levels, the more shielding that occurs. These two factors, nuclear charge
and shielding, are used to explain the trends in atomic radius, ionization energy, and electronegativity.
Questions Modern Atomic Theory:
1. What are the relationships between wavelength, frequency, and energy? You may want to write
down any formulas to support your explanation.
2. Who were the key players in the development of the modern periodic table/modern model of the
atom and how did they alter it? Give detail about each person and about the quantum mechanical
model.
3. Write what you need to remind you about electron configurations and orbital diagrams. Don’t
forget the special cases like the copper family and the chromium family. Use the below terms to
describe or just define them.
a. Ground state
b. Excited state
c. Heisenberg uncertainty principle
d. Probability distribution
e. Pauli exclusion principle
f. Aufbau principle
g. Hund’s rule
4. Describe the four periodic trends; definitions and the trends that they follow in the periodic
table.
5. What is the difference between atomic radius and ionic radius? How does the ionic radius of a
cation compare to the atom from which it was formed? How does the ionic radius of an anion
compare to the atom from which it was formed?
6. What are the special properties of the alkali metals?
TB page 337 Atomic Structure Problems: # 39, 40, 45, 52, 55, 67, 68, 69
AP Chemistry: Electron Configurations, Quantum, and More!
1.
2.
How many protons, neutrons, and electrons are present in each of the following?
a)24 Mg
c)79 X (what element?)
b) 59 Co 3+
d) 59 X 2+ (what element?)
Name the four energy sublevels in order of increasing energy: ____ ____ ____ ____
3. How many orbitals are found in each sublevel? ____ ____ ____ ____
4. Sketch the shapes of the s and p orbitals:
5. Compare and contrast the three 2p orbitals.
6. Explain each of the following:
a) Pauli’s Exclusion Principle b)
Aufbau Principle
c) Hund’s Rule
7. Give the complete electron configuration:
a) fluorine ______________________________________________________________
b) calcium ______________________________________________________________
c) lead
_______________________________________________________________
8. Give the abbreviated electron configuration:
a) sulfur
__________________________________
b) iron
__________________________________
c) iodine
__________________________________
9. How many unpaired electrons are present in the ground state electron configuration: Draw the orbital diagram
for this!
a) aluminum
_____
b) oxygen
_____
c) nickel
_____
10. Identify the following elements.
a) An excited state of this element has the electron configuration: 1s 2 2s2 2p5 3s1
2
4
b) The ground-state electron configuration is [Ne] 3s 3p
_____
_____
c) The ground-state electron configuration contains three unpaired 6p electrons.
_____
11. Write the abbreviated electron configuration:
a) The smallest halogen
________________________
b) The alkali metal with only 2p and 3p electrons
________________________
c) The group 13 element in the same period as tin
________________________
d) One of the nonmetallic elements in group 14
________________________
e) The (yet undiscovered) noble gas after radon
________________________
2
12. A certain oxygen atom has the electron configuration: 1s 2s2 2px1 2py12py13s1
a) How many unpaired electrons are present? __________
b) Is this an excited state of oxygen? __________
c) In going from this state to the ground state, would energy be released or absorbed?
13. From the 1980 AP exam…
a) Write the ground state electron configuration for an arsenic atom, showing the number of electrons in each
subshell.
b) Give one permissible set of four quantum numbers for each of the outermost electrons in a single arsenic
atom when in its ground state.
c) Is an isolated arsenic atom in the ground state paramagnetic or diamagnetic? Explain briefly.
d) Explain how the electron configuration of the arsenic atom in the ground state is consistent with the
existence of the following known compounds: Na3As, AsCl3, and AsF5.
AP Chemistry: Periodic Trends
1. Circle the atom or ion that has the larger atomic radius in each pair.
a) Ca
or
Mg
f) S2- or
S
2+
b) N
or
Be
g) Mg
or
Ca2+
c) Cu
or
Cl
h) F- or
Brd) O
or
Cl
i) Cl- or
Ca2+
e) Na+ or
Na
j) S2- or
Cl
2) Circle the atom or ion with the higher ionization energy in each pair.
a) Ca
or
Br
b) Cl
or
F
c) Mg
or
Al
+
d) Na
or
N
3. What is the difference between a positive electron affinity and a negative electron affinity?
Site specific examples in your explanation.
4. How do the values of ionization energy and electron affinity relate to the electronegativity
of an element?
5. How can the values of ionization energy or electron affinity be used to predict and compare
the strengths of oxidizing or reducing agents?
MULTIPLE CHOICE: CIRCLE THE BEST CHOICE!!
6. In which pair of elements is the larger atom listed first?
A) K, Ca
B) Na, K
C) Cl, S
D) Mg, Na
E) O, N
7. Which of the following is expected to have the largest third ionization energy?
A) Be
B) B
C) C
D) N
E) Al
8. Which one of the following properties of the halogens increases with increasing atomic
mass?
A) ionization energy
B) boiling point of the element
C) tendency to undergo reduction
D) electronegativity
9. Which of the following series is arranged in order of increasing value?
A) The first ionization energies of : oxygen, fluorine, neon
B) The radii of: H – ion, H atom, H + ion
C) The electronegativities of: chlorine, bromine, iodine
D) The boiling points of: iodine, bromine, chlorine
10. Which of the following is most important in determining the periodic trends across a period?
a. Nuclear charge
b. Shielding
c. Increasing numbers of electrons
d. Increasing energy levels
11. Which of the following is (are) important in determining the trend going down a group?
a. Nuclear charge
b. Shielding
c. Increasing numbers of electrons
d. Increasing energy levels
12. Ionization energy is the energy required to remove an electron from a gaseous atom.
Electronegativity is an atom’s ability to pull electrons toward itself in a bond and determines
whether or not a particular bond is ionic, polar or non-polar. Atomic radius is the distance
from the nucleus of the atom to the outermost energy level. Based on your answers to the
previous questions, give the trends going across and down the periodic table for each of
these.
Trend from Left to Right
Trend Down
Ionization
energy
Atomic radii
Electronegativity
13. Put the following in order of INCREASING atomic radius
Arsenic, Potassium, Phosphorus, Chlorine, Titanium
14. Put in order of increasing electronegativity:
P, Co, Zn, Rb, O
15. Which of the following has a greater ionization energy?
a. Nitrogen or Phosphorus?
b. Sodium or Magnesium?
c. Bromine or Hydrogen?
16. When an atom gains electrons to form a negative ion, the increased repulsion between the
electrons causes the radius of the atom to increase. When an atom loses electrons the
decreased repulsion between electrons due to the loss of one causes the radius to decrease.
What happens to the atomic radii when
a. An anion forms?
b. A cation forms?
c. Which atom would be larger?
i. Al3+ or Mg2+?
ii. Cl-1 or S2-?
17. You could use the activity series to figure out the following questions, but don’t do that.
Think about the periodic trends. You will have these questions on your exam when you don’t
have access to a data book.
Based on your knowledge of the trend in ionization energy going down a group
a. which halogen most easily gains electrons?
b. Which halogen most easily loses electrons?
c. Do halogens prefer to lose or gain electrons?
d. Given your answers to and c, predict the trend in reactivity from fluorine to iodine
e. Based on your answers above, predict whether the following reactions will occur
i. Br2 + 2I-1  I2 + 2Br-1
ii. Br2 + 2Cl-1  Cl2 + 2Br-1
AP Chemistry: Atomic Structure Free Response Questions:
1) 1981: 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 wavelengths rather than a continuum of wavelengths?
b. Account for the existence of several series of lines in the spectrum. What quantity
distinguishes one series of lines from another?
c. Draw an electronic energy level diagram for the hydrogen atom and indicate on it the
transition corresponding to the line of the lowest frequency in the Balmer series.
d. What is the difference between an emission spectrum and an absorption spectrum?
Explain why the absorption spectrum of atomic hydrogen at room temperature has only
the lines of the Lyman series.
2) 1982: The values of the first three ionization energies (I1, I2, I3) for magnesium and argon are as
follows:
Mg
Ar
a)
b)
c)
d)
I1 (kJ / mol)
I2 (kJ / mol)
I3
(kJ / mol)
735
1443
7730
1525
2665
3945
Give the electronic configurations of Mg and Ar.
In terms of these configurations, explain why the values of the first and second ionization
energies of Mg are significantly lower than the values for Ar, whereas the third ionization
energy of Mg is much larger than the third ionization energy of Ar.
If a sample of Ar in one container and a sample of Mg in another container are each heated and
chlorine is passed into each container, what compounds, if any, will be formed? Explain in terms
of the electronic configurations given in part (a).
Element Q has the following first three ionization energies:
I1 (kJ / mol)
I2 (kJ / mol)
I3
(kJ / mol)
Q
496
4568
6920
What is the formula for the most likely compound of element Q with chlorine? Explain the choice of
formula on the basis of the ionization energies.
3) 1985: Properties of the chemical elements often show regular variations with respect to their
positions in the periodic table.
a. Describe the general trend in acid-base character of the oxides of the elements in the third
period (Na to Ar). Give examples of one acidic oxide and one basic oxide and show with
equations how these oxides react with water.
b. How does the oxidizing strength of the halogen element vary down a group? Account for
this trend.
c. How does the reducing strength of the alkali metals vary down the group? Account for
this trend.
4) 1987: Use the details of modern atomic theory to explain each of the following observations:
a. Within a family such as the alkali metals, the ionic radius increases as the atomic number
increases.
b. The radius of the chlorine atom (0.99 angstroms) is smaller than the radius of the chloride
ion (1.81 angstroms).
c. The first ionization energy of aluminum (577 kJ/mol) is lower than the first ionization
energy of magnesium (738 kJ/mol)
d. For magnesium, the difference between the 2nd and 3rd ionization energies is much larger
than the difference between the 1st and 2nd energies. (Ionization energies for Mg: 1st
738kJ/mol ; 2nd 1450kJ/mol; 3rd = 7732 kJ/mol).
5)
1990: The diagram shows the first ionization energies from Li to Ne. Briefly (in one to three
sentences) explain each of the following in terms of atomic structure.
a. In general, there is an increase in the first ionization energy from Li to Ne.
b. The first ionization energy of B is lower than that of Be.
c. The first ionization energy of O is lower than that of N.
d. Predict how the first ionization energy of Na compares to those of Li and of Ne. Explain.
6) 1994: Use principles of atomic structure and/or chemical bonding to answer each of the
following:
a. The radium of the Ca atom is 0.197 nanometers; the radius of the Ca2+ ion is 0.099
nanometers. Account for this difference.
b. The lattice energy of CaO(s) is -3,460 kJ per mole; the lattice energy of K2O (s) is -2,240
kJ per mole. Account for the difference.
c. Explain the difference between Ca and K in regard to (i) their first ionization energies (ii)
their second ionization energies.
Atom first ionization energy (kJ per
second ionization energy (kJ per mole)
mole)
K
419
3050
Ca
590
1140
d. The first ionization energy of Mg is 738 kJ per mole and that of Al is 578 kJ per mole.
Account for this difference.
7) 1970: Explain why in aqueous solution,
(a)
Ti3+ is colored but Sc3+ is not.
(b)
Ti2+ is a reducing agent but Ca2+ is not.
8) 1970: What is meant by the lanthanide contraction? Account for this phenomenon. Give two
examples of its consequences.
9) 1971: There is a greater variation between the properties (both chemical and physical) of the
first and second of a group or family in the periodic table than between the properties of the
second and third members of the group. Consider as examples either the group containing
nitrogen or the one containing oxygen. Select three properties and discuss the variation of
these properties to illustrate the generalization expressed in the first sentence of the question.
10) 1972: Consider the following melting points in degrees Celsius:
Alkali metals
Halogens
Li 181˚
F2 –119˚
Na 98˚
Cl2 –101˚
K
63˚
Br2
–7˚
Rb 39˚
I2 +104˚
Cs 29˚
(a)
Account for the trend in the melting points of the alkali metals.
(b)
Account for the trend in the melting points of the halogens.
(Make sure that your discussion clarifies the difference between the two trends.)
11) 1973 D:
First ionization Energy
(kilocalories/mole) Covalent Radii, Å
Li
124 1.34
Be
215 0.90
B
191 0.82
C
260 0.77
N
336 0.75
O
314 0.73
F
402 0.72
The covalent radii decrease regularly from Li to F, whereas the first ionization energies do not.
For the ionization energies, show how currently accepted theoretical concepts can be used to
explain the general trend and the two discontinuities.
12) 1976 D: M(s) + Cl2(g)  MCl2(s)
The reaction of a metal with chlorine proceeds as indicated above. Indicate, with reasons for
your answers, the effect of the following factors on the heat of reaction for this reaction.
(a)
A large radius versus a small radius for M2+
(b)
A high ionization energy versus a low ionization energy for M.
13) 1977 D
The electron affinities of five elements are given below.
13Al 12 kcal/mole
14Si 32 kcal/mole
15P 17 kcal/mole
16S 48 kcal/mole
17Cl 87 kcal/mole
Define the term “electron affinity” of an atom. For the elements listed above, explain the
observed trend with the increase in atomic number. Account for the discontinuity that occurs at
phosphorus.
14) 1984 D: Discuss some differences in physical and chemical properties of metals and nonmetals.
What characteristic of the electronic configuration of atoms distinguishes metals from
nonmetals. On the basis of this characteristic explain why there are many more metals than
nonmetals.
*** Separate paper:
Atomic Structure Multiple Choice
PES Notes and Video
PES Problem Set
AP Chemistry: Structure and Periodicity Review
1.
Complete the chart:
P+
n0
a.
Sn-120
33 2b.
S
e-
Electron configuration
16
c.
d.
e.
f.
g.
Mo
Cu2+
W
N3Ag
2. Calculate the energy in Joules/photon and in KJ/mol for red light with a wavelength of 665.7nm.
3. A laser emits a pulse at 450nm that contains 2.75 x 10 15 photons what is the total energy of the pulse.
4. For the following sets of quantum numbers, determine which orbitals are nonexistent:
n
l
ml
s
a.
5
3
-2
-1
b.
5
3
-3
½
c.
3
3
-3
½
d.
3
0
0
-½
e.
2
1
1
f.
1
0
-1
g.
4
2
-2
h.
3
3
0
5. What are the possible quantum numbers for a 3d electron?
6. How many valence electrons does each of the following elements have? P, Cr, Kr, I, Ba, O, Cd.
7. For nitrogen, how many different excited-state configurations are there in which no electron has n > 2?
Write them.
8. Arrange the following oxides in increasing order of acidity: CO 2, CaO, Al2O3, SO3, SiO2, P2O5
9. List the correct set of quantum numbers for the valence electrons of carbon and calcium.
10. Arrange the following atoms in order of increasing size: Cl, F, P, S
11. Element X has the following values of for ionization energy. Name an element that would have similar
properties to element X. Predict the formula for an oxide of X.
IE1
IE2
IE3
503
965
3600
12. Use periodic trend to explain why 4A elements with higher Z#’s tend to form +4 ions and the lower Z#’s
tend to form -4.
13. Arrange the following in order of decreasing ionization energy: Br, Ar, Ar +, and Cl
14. Calculate the energy change associated with the following electron transitions. Indicate the direction of
the energy. Calculate the wavelength and predict where the photon fall in the electromagnetic spectrum.
a. n=3 to n=5
b. n=4 to n=2
Don’t forget the history stuff.
AP CHEMISTRY: THE
EMISSION SPECTRUM OF
HYDROGEN ATOM
INTRODUCTION:
In the early 1900s, Ernest Rutherford devised an experiment which revised the model of the
atom. His gold foil experiment showed that most of the mass (including the protons) is found in the
nucleus and that the electrons were in a region about that nucleus.
Line spectra for elements would later be used to refine the model of the atom. The main question
to be answered was why atoms emitted light of only certain frequencies instead of a continuous
spectrum like that of an incandescent light bulb. Niels Bohr answered this question and introduced the
ideas of quantization and energy levels.
In your chemical studies, you learned that after electrons are excited from the ground state (by
an external source of energy), they naturally fall back to lower energy levels emitting energy in the
process. It has been determined that electrons that fall to the first energy level emit energy in the
range of ultraviolet light [Lyman Series], visible light is emitted if they fall to the second energy level
[Balmer Series], and infrared light emitted if they fall to the third energy level [Paschen Series].
In this experiment, you will use the equations developed by Balmer, Rydberg, and Bohr to predict
the wavelengths and energies of light emitted for transitions in the hydrogen atom. These values will be
compared with your experimental values and the accepted values found in your text.
Purpose:
The purpose of this experiment is to measure the wavelengths and calculate the energies of the
lines in the Balmer series of the hydrogen atom. These values will be compared with the values
predicted from the equations.
Equipment/Materials:
hydrogen discharge tube
spectroscope
Safety: This experiment poses no unusual safety hazards. Do not touch the discharge tubes because
they can become quite hot.
Procedure:
1. Point the spectroscope outside the classroom and record the visible spectrum that you observe.
Label the wavelength and colors that you see. You should be able to calculate the wavelength from
the formulas here= The numbers on the spectroscope scale represent 1,000 Angstroms (Ex. 5.20 =
5200 angstroms). 1.0x 1010 angstroms =1 meter.
2.
Point the spectroscope to the hydrogen light source. Measure and record the wavelengths of the
four lines in the Balmer series of the hydrogen atom. You will have difficulty measuring one of the
lines close to 400 nm, so make careful observations.
3. Using the Rydberg equation, calculate the wavelength that would be predicted for each transition in
the Balmer series. All transition are from higher energy levels to the n = 2 energy level.
A little History with your calculations: These transitions are named after J.J. Balmer because he was
the first to describe the lines according to the following equation:
1 
 1
 R 2  2 

n 
2
1
Where n =3,4,5… and R = Rydberg constant of -2.178 x 10-18 J
Balmer did not know at the time that this equation represented the electrons transitioning
from one energy level to another. It was Bohr who combined the work of Rydberg and
Planck to come up with his model and changed Balmer’s equation slightly to:
 1
1
E  R '  2  2
n
 lower nupper



where nlower is the ground state energy level and nupper is the excited state energy level and
R = -2.178x10-18 J.
(Since you probably did not get a whole number in number 3, compare to the nearest whole number that
makes sense to calculate your percent error for each transition.)
Predicted Wavelengths (use Rydberg’s equation; show work):
n=6
n=5
n=4
n=3
Summary of
to
to
to
to
n
n
n
n
=2
=2
=2
=2
Observed wavelength, nm
results
Calculated wavelength,
Accepted wavelength, nm
nm
(use different references)
Line 1
Line 2
Line 4
Line 5
4. Compare your experimental results and provide reasons why your results differ from the calculated
and “accepted” values?
5. Construct an energy level diagram showing the energies and colors of light observed in the Balmer
series.
6. When you “excited” the hydrogen gas, you obtained an emission spectrum. You can also shine white
light into a hydrogen gas producing an absorption spectrum. What is the difference between an
emission spectrum and an absorption spectrum? Draw an absorption spectrum for hydrogen.
7. Explain why you obtained an emission spectrum for hydrogen that only showed distinct lines
(wavelength) while you obtained an emission spectrum showing bands of colors (all wavelength) when
you looked at the emission spectrum from sunlight?
Analysis Question:
1. The Paschen series lines in the atomic spectrum of hydrogen result from transitions from n > 3 to
n = 3. Calculate the wavelength in nm of a line in this series resulting from an n = 6 to n = 3 transition.
Where would this line be found in the electromagnetic spectrum?
2. Calculate the wavelength of light emitted when an electron changes from n = 3 to n = 1 in the hydrogen
atom. In what region of the spectrum is this radiation found?
3. a) What is the energy of an electron that is at ground state (lowest level, n=1)?
b) How much energy must be added to the hydrogen atom in order to cause the electron to move to the
principle quantum, n = 5?
4. How much energy is released from a hydrogen atom when an electron moves from n=6 to n=2?
b) What is the wavelength of the quantum of light released? What color will it be?
Discuss possible sources for that error AND explain how those sources of error affected the
calculations. Please do not discuss “calculation errors,” “human errors,” or “measurement errors” because
these are not specific enough for error analysis.
AP Chemistry
Flame Test Lab
Introduction:
For most Americans, a Fourth of July celebration is not complete without spectacular fireworks and
the crowd’s “oohs” and “ahhs”. Did you ever wonder what caused the brilliant array of firework
colors? These colors are usually the result of the interaction of an ionic compound with the
extreme temperatures present in the firework’s explosion.
By placing atoms of a metal into a flame, electrons can be induced to absorb energy and jump to an
excited energy state, a quantum jump. They then return to their ground state by emitting a photon
of light (the law of conservation of energy indicates that the photon emitted will contain the same
amount of energy as that absorbed in the quantum jump). The amount of energy in the photon
determines its color; red for the lowest energy visible light, increasing energy through the rainbow
of orange yellow green blue indigo, and finally violet for the highest energy visible light. Photons
outside the visible spectrum may also be emitted, but we cannot see them.
The arrangement of electrons in an atom determines the sizes of the quantum jumps, and thus the
energy and colors of the collection of photons emitted, known as an emission spectrum. In this way
the emission spectrum serves as a ‘fingerprint’ of the element to which the atoms belong. We can
view the emission spectrum of colors all at once with the naked eye. It will appear to be one color,
which we will carefully describe. It is also possible to view the separate colors of the emission
spectrum by using a diffraction grating or spectroscope. These instruments bend light of different
energies differently. Low energy red light is bent the most, and high energy violet the least. This
allows us to see the various distinct colors of the emission spectrum of a sample.
In this lab, you will record the flame test color of several substances by placing droplets of
solutions of ionic salts onto a nichrome wire loop, then placing the loop into a bunsen burner flame.
Next, you will attempt to use a spectroscope to view the separate colors of the emission spectra,
but this is difficult to do under our lab conditions because the flame test is of short duration and
the lab lighting can only be adjusted to a certain degree.
Cobalt blue glass filters are often used when viewing mixtures of metals to screen out light that is
yellow in color. The human eye sees yellow very well, since it is in the middle of the spectrum visible
to the eye. Colors at the edges of the visible spectrum, especially violet, are more difficult to see.
Cobalt glass absorbs light in the yellow wavelengths, but is transparent to light of higher energy
(this is why it looks blue!). Viewing a yellow flame through cobalt glass will allow us to see if there is
any higher energy light present.
Equipment:
Bunsen burner
1.0M HCl
Loops of nichrome wire
Various salt solutions in dropper bottles
Cobalt blue glass
Diffraction gratings or spectroscopes
Procedure:
Each lab bench has one or two aqueous salt solutions on it. You will travel from station to
station, testing every solution in the Bunsen burner flame. To test a solution, first clean the
nichrome wire by dipping it into 1.0M HCl and then heating it in the hottest part of the burner flame
until there is a constant yellow color. (This yellow flame is NOT emission by the nichrome wire; it is
a result of the hot Bunsen burner flame being cooled by the wire.) Do this twice to ensure removal
of any contaminants. Add a single drop of the salt solution to the loop at the end of the wire. Place
this drop of solution into the hottest part of the burner flame.
After recording the observed color, look at the colored flame through a spectroscope.
Describe and/or illustrate the observed spectrum in your data table. If possible, estimate and
record numerical values for the brightest emission lines you see. Finally, for specified solutions,
look at the flame through cobalt blue glass.
Data:
Create a data table on a separate sheet of paper. For each substance, you need to record two
things: the color of the flame and an illustration of the emission spectrum (with labels of bright
line spectra and approx. wavelengths or colors).
Discussion Questions:
1. What is the difference between a continuous emission spectrum and a line emission spectrum?
Which one did you observe from the solutions in this experiment?
2. List the colors of the visible spectrum in order of increasing wavelength. Don’t look in the book!
Try it from memory!
3. Define frequency. What are the units of frequency? Give the relationship between frequency
and wavelength.
4. Qualitatively describe the relationship between frequency and energy.
5. Several compounds contain the same cation (NaCl and Na2SO4; KNO3 and K2SO4; etc.) and
several contain the same anion (NaCl, CaCl2, SrCl2, etc.). If the cation is responsible for the
flame test color, what flame test result would you expect to see from compounds containing the
same cation? Compare the flame test colors for a set of these compounds. Do your data
suggest that the cation or the anion is responsible for the flame test color? How do you know?
6. Define ground state and excited state.
7. Describe the activity of electrons when a substance is vaporized in a flame. Explain what is
viewed through a spectroscope or diffraction grating and how such an instrument helps in
identifying an unknown substance.
8. (a) The frequency of a red spectral line is 1.60 x 1014 Hz. How much energy does each photon of
this light have?
(b) The frequency of a green spectral line is 2.50 x 1014 Hz. How much energy does each photon
of this light have?
(c) Considering your calculations from a and b above, explain why ultraviolet light can cause
cancer in high doses, while infrared radiation is harmless. Refer back to your answers to a
and b in this response.