Electrons, Energy, & the Electromagnetic
Spectrum Notes
Simplified, 2D Bohr Model:

Here’s how the type/form of EM radiation can be determined…
The
amount of energy released when an electron falls from a higher to a lower
energy level is directly proportional to its frequency. The calculation follows the equation: E = h ^{.} ν
E
= Energy (unit is J)
h = Planck’s Constant (6.626 x 10^{34} J^{.}s)
ν = frequency (unit is Hz or ^{1}/_{second})
EXAMPLE
1: A particle of EM radiation has an
energy of 1.15 x 10^{16} J.
What is its frequency?
1.15
x 10^{16} J = 6.626 x 10^{34} J^{.}s ^{.}
ν
ν = 1.74 x 10^{17} Hz
The
type of electromagnetic radiation can be determined if one knows the
wavelength. The wavelength is inversely
proportional to the frequency. The
calculation follows the equation: c =
λ ^{.} ν
c = speed of light (3.00 x 10^{8} m/s)
λ = wavelength (unit is m)
ν = frequency (unit is Hz or ^{1}/_{s})
EXAMPLE 2: What is the wavelength of the
same particle from EXAMPLE 1?
3.00
x 10^{8} m/s = λ ^{.} 1.74 x 10^{17} Hz
λ
= 1.72 x 10^{9} m
EXAMPLE
3: What type of electromagnetic
radiation is the particle from EXAMPLE 1?
answer for wavelength is 10^{9} so use chart on the next page to determine…
xrays or ultraviolet
(either one is acceptable)
PROBLEMS FOR YOU TO TRY ON YOUR OWN…
1.)
A particle of EM radiation has a frequency of 5.76 x 10^{14} Hz.
(A) How much energy does this
particle have?
(B) What is the wavelength of
this particle?
(C)
What specific type of electromagnetic radiation does this particle represent?
2.)
A particle of electromagnetic radiation has 2.39 x 10^{13} Joules of
energy.
(A) What is the wavelength of
this particle?
(B)
What type of electromagnetic radiation does this particle represent?
Light Calculations Notes:
* Frequency and wavelength are ___________________ proportional
* Energy and frequency are ____________________ proportional
Light as a Particle Notes:
* Object emits energy in small, specific amounts (called ________________)
* _________________: particle of EM radiation carrying a quantum of energy
* Einstein suggested that light had properties of both waves and particles
Referred to as the ________________________________________ of
light
Quantum Theory Notes:
* When atom falls from excited state to ground state,
____________________________________________
_______________________________________________________________________________________
* Energy of photon = difference
______________________________________________________________
_______________________________________________________________________________________
* Energy states of atoms are fixed
Bohr model of the hydrogen atom Notes:
* said that e circled the nucleus in fixed paths
* when in path, has fixed amount of energy
* e cannot exist in space between path
* drawback of Bohr’s model =
EMISSION & ABSORPTION SPECTRA NOTES
According to the Bohr atomic model, electrons
orbit the nucleus within specific energy levels. These levels are
defined by unique amounts of energy. Electrons possessing the lowest
energy are found in the levels closest to the nucleus. Electrons with
higher energy are located in progressively more distant energy levels.
If an electron absorbs sufficient energy to bridge
the "gap" between energy levels, the electron may jump to a higher
level. Since this change results in a vacant lower orbital, this
configuration is unstable. The "excited" electron releases its
newly acquired energy as it falls back to its initial or ground state.
Often, the excited electrons acquire sufficient energy to make several energy
level transitions. When these electrons return to the ground state,
several distinct energy emissions occur. The energy that the electrons
absorb is often of a thermal or electrical nature, and the energy that an
electron emits when returning to the ground state is called electromagnetic
radiation.
In 1900, Max Planck studied visible
emissions from hot glowing solids. He proposed that light was emitted in
"packets" of energy called quanta, and that the energy of each
packet was proportional to the frequency of the light wave. According to
Einstein and Planck, the energy of the packet could be expressed as the product
of the frequency (n) of emitted light and Planck's constant (h). The equation is
written as E = hn
If white light passes through a prism or
diffraction grating, its component wavelengths are bent at different
angles. This process produces a rainbow of distinct colors known as a continuous
spectrum. If, however, the light emitted from hot gases or energized
ions is viewed in a similar manner, isolated bands of color are observed.
These bands form characteristic patterns  unique to each element. They
are known as bright line spectra or emission spectra.
By analyzing the emission spectrum of hydrogen
gas, Bohr was able to calculate the energy content of the major electron
levels. Although the electron structure as suggested by his planetary
model has been modified according to modern quantum theory, his description and
analysis of spectral emission lines are still valid.
In addition to the fundamental role of spectroscopy
played in the development of today's atomic model, this technique can also be
used in the identification of elements. Since the atoms of each element
contain unique arrangements of electrons, emission lines can be used as
spectral fingerprints. Even without a spectroscope, this type of
identification is possible since the major spectral lines will alter the color.
Properties of Light Worksheet
Part 1  Select the best answer
1. Which has a longer wavelength, orange or
violet light?
2. Which has a higher energy, xrays or gamma rays?
3. Which has a lower frequency, radio waves or green light?
4. Which has the shortest wavelength, violet or ultraviolet light?
5. Which has lower energy, infrared light or xrays?
Part 2  Fill in the blanks
6. _______________ formed a theory to
explain the structure of an atom by revising physical theories.
7. As the energy level increases, the amount of energy an electron will
possess _______________.
8. Electrons give off energy in finite amounts called _______________
when returning to the ground state.
9. When this energy is released in the form of light it is called a
_______________.
10. The speed of light = _______________ (give number and units)
11. The symbol for wavelength is _______________.
12. In the equation c = λ ^{.} ν, c represents _______________, ν
represents _______________, and
λ
represents _______________.
13. In the equation c = λ ^{.} ν, λ and ν are _______________ proportional.
14. In the equation E = h ^{.} ν, h
represents _______________ and E represents _______________.
15. In the equation E = h ^{.} ν, E
and ν are _______________ proportional.
16. Bohr chose the element _______________ to prove his theory.
Part 3  True or False
17. Electrons may regularly occupy spaces between
orbitals.
18. The varying wavelengths on the electromagnetic radiation spectrum
travel at different speeds.
19. Atoms release energy when electrons jump to higher energy levels.
EM SPECTRUM, WAVELENGTH, FREQUENCY, AND ENERGY WORKSHEET
1.) Look at the EM spectrum below to answer this
question.
As you
move across the visible light spectrum from red
to violet…
(A) Does the wavelength increase or decrease?
(B) Does the frequency increase or decrease?
(C) Does the energy increase or decrease?
2.) A beam of microwaves has a frequency of 1.0 x
10^{9} Hz. A radar beam has a
frequency of 5.0 x 10^{11} Hz.
Which type (microwave or radar)…
(A) has a longer wavelength?
(B) is closer to visible light on the EM spectrum?
(C) is closer to xrays in frequency value?
3.) What is the frequency of an EM radiation wave
if its wavelength is 3.6 x 10^{9} meters?
4.) A beam of EM radiation has a wavelength of
4.257 x 10^{7} cm. What is its
frequency?
5.) A photon of light has a wavelength of 3.20 x
10^{5} meters. Find…
(A)
the frequency
(B)
the energy
(C)
the region of the EM spectrum/type of radiation
6.) A photon has an energy of 4.00 x 10^{19}
J. Find…
(A)
the frequency
(B)
the wavelength
(C)
the region of the EM spectrum/type of radiation
7.) A bright line spectrum contains a line with a
wavelength of 518 nm. Determine…
(A)
the wavelength in meters
(B)
the frequency
(C)
the energy
(D)
the color
*8.) Cobalt60 is an artificial radioisotope that is produced in a nuclear
reactor for use as a gamma ray source in the treatment of certain types of
cancer. If the wavelength of the gamma
radiation from a cobalt60 source is 1.00 x 10^{3} nm, calculate the
energy of a photon of this radiation.
ELECTRON ARRANGEMENT NOTES
Heisenberg
Uncertainty Principle:
ELECTRON ARRANGEMENT
WORKSHEET
1. What is
an electron cloud?
2. Name the
three major divisions within an electron cloud with respect to the energy of an
electron.
3. What
letter represents the principal quantum number?
4. What
does the principle quantum number tell about an electron?
5. What
formula is used to determine the maximum number of electrons that can occupy
any energy level?
6. What is
the maximum number of electrons for each of the following?
(A)
1st energy level (B) 4th
energy level (C) n = 3 (D) n = 5
7. Energy
levels are divided into _______________.
8. How can
we determine the possible number of sublevels in any energy level?
9. Name the
four primary sublevels in order of increasing energy.
10. Circle
the sublevel that represents the lowest
energy in each pair.
(A)
1s or 2s (B) 2s or 2p (C) 4f or 4d (D) 3d or 4s (E) 7s or 5d
(F)
6s or 4s (G) 4p or 5p (H) 3s or 3d (I) 2p or 3s
11. Sublevels are divided into _______________.
12. Each
orbital can hold up to _______________ electrons.
13. Sketch
the shapes of the orbitals for the sublevels listed.
(A)
s: (B) p_{x}:
(C) p_{y}: (D) p_{z}:
14. How many orbitals are in each sublevel?
(A)
s _______________ (B) p _______________
(C) d _______________ (D) f _______________
More
Electron Arrangement Notes!
Examples:
Se 1s^{2} 2s^{2}
2p^{6} 3s^{2} 3p^{6} 4s^{2} 3d^{10} 4p^{4}
Sn 1s^{2} 2s^{2}
2p^{6} 3s^{2} 3p^{6} 4s^{2} 3d^{10} 4p^{6}
5s^{2} 4d^{10} 5p^{2}
Hg 1s^{2}
2s^{2} 2p^{6} 3s^{2} 3p^{6} 4s^{2} 3d^{10}
4p^{6} 5s^{2} 4d^{10} 5p^{6} 6s^{2} 4f^{14}
5d^{10}
HOEL
(Highest Occupied Energy Level): energy level
furthest from the nucleus that contains at least one electron
How to
determine this using electron configuration?
~ largest nonexponent number
Se 1s^{2} 2s^{2}
2p^{6} 3s^{2} 3p^{6} 4s^{2} 3d^{10}
4p^{4} HOEL
= 4
Sn 1s^{2} 2s^{2}
2p^{6} 3s^{2} 3p^{6} 4s^{2} 3d^{10} 4p^{6}
5s^{2} 4d^{10} 5p^{2} HOEL = 5
Hg 1s^{2}
2s^{2} 2p^{6} 3s^{2} 3p^{6} 4s^{2} 3d^{10}
4p^{6} 5s^{2} 4d^{10} 5p^{6} 6s^{2}
4f^{14} 5d^{10} HOEL
= 6
Valence
Electrons: electrons in the HOEL
How to
determine this using electron configuration?
~ add up exponents of terms in HOEL
Se 1s^{2} 2s^{2}
2p^{6} 3s^{2} 3p^{6} 4s^{2} 3d^{10} 4p^{4}
HOEL
= 4 Valence electrons =
2 + 4 = 6
Sn 1s^{2} 2s^{2}
2p^{6} 3s^{2} 3p^{6} 4s^{2} 3d^{10} 4p^{6}
5s^{2} 4d^{10}
5p^{2}
HOEL
= 5 Valence electrons =
2 + 2 = 4
Hg 1s^{2}
2s^{2} 2p^{6} 3s^{2} 3p^{6} 4s^{2} 3d^{10}
4p^{6} 5s^{2} 4d^{10} 5p^{6} 6s^{2} 4f^{14} 5d^{10}
HOEL
= 6 Valence electrons =
2
Noble
Gas Configuration: shortcut for electron
configuration
How is it
written?
~ [ symbol for noble has closest to
element with lower atomic # ]
~ [after brackets] next number is
the period that the element is located in
~ after that number, write “s”
~ continue electron configuration in
diagonal rule order until appropriate # of
electrons is reached
*NOTE: ending of electron configuration and noble
gas configuration should be the same*
Se 1s^{2} 2s^{2}
2p^{6} 3s^{2} 3p^{6} 4s^{2} 3d^{10} 4p^{4}
[Ar] 4s^{2}
3d^{10} 4p^{4}
^{18 20 30 34}
Sn 1s^{2} 2s^{2}
2p^{6} 3s^{2} 3p^{6} 4s^{2} 3d^{10} 4p^{6}
5s^{2} 4d^{10} 5p^{2}
[Kr] 5s^{2} 4d^{10} 5p^{2}
^{ 36 38 48 50 }
Hg 1s^{2} 2s^{2}
2p^{6} 3s^{2} 3p^{6} 4s^{2} 3d^{10} 4p^{6}
5s^{2} 4d^{10} 5p^{6} 6s^{2} 4f^{14} 5d^{10}
[Xe]
6s^{2} 4f^{14} 5d^{10}
^{ 54 56 70 80}
Orbital Notation: drawing of how
electrons are arranged in orbitals; will only need to do this for the HOEL
*NOTE: ___
= orbital
or =
electrons
Se 1s^{2} 2s^{2}
2p^{6} 3s^{2} 3p^{6} 4s^{2} 3d^{10} 4p^{4}
Sn 1s^{2} 2s^{2}
2p^{6} 3s^{2} 3p^{6} 4s^{2} 3d^{10} 4p^{6}
5s^{2} 4d^{10} 5p^{2}
Hg 1s^{2} 2s^{2}
2p^{6} 3s^{2} 3p^{6} 4s^{2} 3d^{10} 4p^{6}
5s^{2} 4d^{10} 5p^{6} 6s^{2} 4f^{14} 5d^{10}

Dot Diagrams: symbol represents nucleus
and nonvalence (“innershell”) electrons;
dots around symbol represent valence electrons
is final answer
ELECTRON DOT DIAGRAMS WORKSHEET

ELEMENT 
ELECTRON
CONFIGURATION 
NOBLE GAS
CONFIGURATION 
HIGHEST OCCUPIED
ENERGY LEVEL 
# OF VALENCE
ELECTRONS 
ORBITAL NOTATION
OF H.O.E.L. 
ELECTRON DOT
DIAGRAM 
1. EX 
magnesium 
1s^{2}
2s^{2} 2p^{6} 3s^{2} 
[Ne]
3s^{2} 
3 
2 

: Mg 
2. 
carbon 





C 
3. 
sulfur 





S 
4. 
barium 





Ba 
5. 
nickel 





Ni 
6. 
oxygen 





O 
7. 
arsenic 





As 
8. 
lead 





Pb 
9. 
lithium 





Li 
10. 
neon 





Ne 
11. 
bromine 





Br 
12. 
sodium 





Na 
13. 
chlorine 





Cl 
14. 
argon 





Ar 
15. 
calcium 





Ca 
16. 
zinc 





Zn 
17. 
potassium 





K 
18. 
iodine 





I 
19. 
cobalt 





Co 
20. 
nitrogen 





N 
21. 
fluorine 





F 
22. 
iron 





Fe 
23. 
phosphorus 





P 
24. 
aluminum 





Al 
QUANTUM NUMBERS NOTES
~ describe one specific electron
~ 1st quantum number = PRINCIPAL QUANTUM NUMBER
~
abbreviated "n"
~
tells the energy level the electron is located in
~
n = number of the energy level
~
1st energy level: n = 1, 4th energy level: n = 4, etc.
~ 2nd quantum number = ANGULAR MOMENTUM QUANTUM NUMBER
~ abbreviated " l "
~ tells the sublevel the
electron is located in
~ tells shape of orbital
~ "s" sublevel: l = 0, "p" sublevel: l = 1,
"d" sublevel: l = 2, "f" sublevel: l = 3
~ 3rd quantum number = MAGNETIC QUANTUM NUMBER
~ abbreviated "m"
~ tells which orbital the
electron is in
~ tells orientation of orbital
around nucleus
~ m =  l
.. + l
~ 4th quantum number = SPIN QUANTUM NUMBER
~ abbreviated " s "
~ tells which electron is
being described
~ tells which way electron is
spinning
~ s = ^{1}/_{2} or +^{1}/_{2}
Pauli Exclusion
Principle:
EXAMPLE QUESTIONS:
1.) What are the 4 quantum numbers for
the electron shown above?
2.) If the electron in question 1 was
the last electron added, what element would it be?
3.) Draw in the electron (and the
orbital notation) for the electron with the following quantum numbers.
n
= 3 l = 2 m = 1 s =  ˝
4.) How many electrons in an atom can
have the quantum numbers n = 3 and l = 1?
5.) What are the four quantum numbers
for the electron circled in the diagram below?
n
= l = m = s =
QUANTUM NUMBERS WORKSHEET

Element 
1s 
2s 
2p 
3s 
3p 
3d 
4s 
4p 
4d 
4f 
5s 
1.) 
K 
_ 
_ 
_ _ _ 
_ 
_ h _ 
_ _ _ _ _ 
_ 
_ _ _ 
_ _ _ _ _ 
_ _ _ _ _ _ _ 
_ 
2.) 
O 
_ 
h 
_ _ _ 
_ 
_ _ _ 
_ _ _ _ _ 
_ 
_ _ _ 
_ _ _ _ _ 
_ _ _ _ _ _ _ 
_ 
3.) 
Ar 
_ 
_ 
_ _ i 
_ 
_ _ _ 
_ _ _ _ _ 
_ 
_ _ _ 
_ _ _ _ _ 
_ _ _ _ _ _ _ 
_ 
4.) 
Br 
_ 
_ 
_ _ _ 
_ 
_ _ _ 
_ _ _ _ _ 
_ 
_ _ _ 
_ _ _ _ _ 
_ _ _ _ _ _ _ 
_ 
5.) 
Rb 
_ 
_ 
_ _ _ 
_ 
_ _ _ 
_ _ _ _ _ 
_ 
_ _ _ 
_ _ _ _ _ 
_ _ _ _ _ _ _ 
_ 
6.) 
Co 
_ 
_ 
_ _ _ 
_ 
_ _ _ 
_ _ _ _ _ 
_ 
_ _ _ 
_ _ _ _ _ 
_ _ _ _ _ _ _ 
_ 
7.) 
Se 
_ 
_ 
_ _ _ 
_ 
_ _ _ 
_ _ _ _ _ 
_ 
_ _ _ 
_ _ _ _ _ 
_ _ _ _ _ _ _ 
_ 
8.) 
B 
_ 
_ 
_ _ _ 
_ 
_ _ _ 
_ _ _ _ _ 
_ 
_ _ _ 
_ _ _ _ _ 
_ _ _ _ _ _ _ 
_ 
9.) 
P 
_ 
_ 
_ _ _ 
_ 
_ _ _ 
_ _ _ _ _ 
_ 
_ _ _ 
_ _ _ _ _ 
_ _ _ _ _ _ _ 
_ 
** Each question below corresponds to the number in the
table. **
For 1 – 3, give the four quantum numbers for the electron indicated.
1.) n = l = m = s
=
2.) n = l = m = s
=
3.) n = l = m = s
=
For 4 – 9, draw in the electron with the following sets of quantum numbers.
4.) n = 3 l = 2 m = 0 s = + ˝ 5.) n = 3 l = 1 m = + 2 s = + ^{1}/_{2}
6.) n = 4 l = 1 m = + 1 s =  ^{1}/_{2}
7.) n = 2 l = 1 m = 0 s = + ^{1}/_{2}
8.) n = 3 l = 2 m =  1 s =  ^{1}/_{2}
9.) n = 2 l = 1 m =  1 s =  ^{1}/_{2}
10.) Does an electron with this set of
quantum numbers exist in the element calcium?
n
= 4 l = 1 m
= 0 s =  ˝
Determine the element whose outermost electron (last electron added) is being
defined by the following quantum numbers.
11.) n = 1 l =0 m = 0 s = ^{1}/_{2}
12.) n = 4 l =1 m = 1 s = ^{1}/_{2}
13.) n = 3 l =1 m = 1 s = +^{1}/_{2}
14.) n = 4 l = 0 m = 0 s = +^{1}/_{2}
15.) n = 3 l = 2 m = 2 s = ^{1}/_{2}
Unit 4 Review
Section I  Problems Given:
E = h ^{.} ν h = 6.626 x 10^{34} J^{.}s
c
= λ ^{.} ν c
= 3.00 x 10^{8} m/s
1. What is the frequency of a wave with a
wavelength of 3.5 x 10^{4} m?
2. What is the energy of a photon with a
frequency of 5.41 x 10^{17} Hz?
3. What type of electromagnetic radiation is
described in question 2?
Section II  Electromagnetic Spectrum
1. Label both ends of the spectrum with high/low
frequency, high/low energy, and long/short wavelength
radio waves microwaves infrared light ROYGBIV ultraviolet
light xrays gamma rays
3. Which has a shorter wavelength, radio or
ultraviolet?
4. Which has a lower frequency, yellow or green
light?
5. In the equation E = h ^{.}
ν, energy and frequency are ____________________
proportional.
6. In the equation c = λ ^{.} ν, wavelength and
frequency are __________________ proportional.
7. The symbol for wavelength is _____.
8. Electrons give off energy in the form of a
____________________ when returning to the ground state.
9. Which scientist proposed the idea that
electrons travel around the nucleus in fixed paths?
10. When an electron moves from the ground state to the excited state, energy is ____________________.
11. Bohr chose the element ____________________ to prove his theory.
12. The dual
waveparticle nature of electrons describes how the electrons in atoms can
behave as
____________________ and
____________________.
Section III  Electrons
1. What is an electron cloud?
2. Who proposed the uncertainty principle?
3. Who is credited with the idea that electrons are placed in the lowest energy level first?
4. What rule
requires that each of the "p" orbitals (at a particular energy level)
receive one electron before any
of the orbitals can have two
electrons?
5. What is the maximum number of electrons in any orbital?
6. The principal quantum number, n, indicates the ____________________.
7. The maximum
number of electrons in an energy level can be determined by the equation
________________
That means the maximum number of electrons in the 3rd
energy level is ____________________.
8. The number of sublevels in any energy level can be determined by ____________________.
9. The number
of orbitals in an energy level can be determined by the equation
____________________.
So, the 3rd energy level has _____ orbitals.
(_____ is/are "s" orbitals, _____ is/are "p" orbitals, and
_____
is/are "d" orbitals)
10. List the four sublevels according to increasing energy.
11. The "s" sublevel is shaped like a ____________________ and has _____ orbitals.
12. A "p" sublevel is shaped like a ____________________ and has _____ orbitals.
13. The "d" sublevel has _____ orbitals and the "f" sublevel has _____ orbitals.
Section IV  Electron configuration, noble gas configuration, valence electrons, orbital notations
1. What is the electron configuration for phosphorus?
2. How many total electrons are in a neutral atom of phosphorus?
3. Write the noble gas configuration for phosphorus.
4. What is the highest occupied energy level for phosphorus?
5. What is the atomic number of phosphorus?
6. Draw the orbital notation for phosphorus.
7. Circle the
last electron added to phosphorus. What
are the four quantum numbers for this electron?
n
= l
= m = s =
8. How many electrons are in the highest occupied energy level of phosphorus?
9. How many innershell electrons does phosphorus have?
10. In which orbitals are the innershell electrons located?
11. Draw the electron dot diagram for phosphorus.
Section V  Quantum numbers (Honors level only)
1. How many electrons can be described by the quantum numbers n = 3 and l = 1?
2. How many electrons in an atom can have the quantum numbers n = 2 and l = 3?
3. How many electrons can have the value n = 3?
4. How many electrons in an atom have the quantum numbers n = 4 and l = 2?
5. Which of
the following sets of quantum numbers does NOT represent a possible set of
quantum numbers?
(There may be more than one correct
answer.)
n l m
s
(A) 4 8 4 ^{1}/_{2}
(B) 6 5 5 ^{1}/_{2}
(C) 3 2 2 ^{1}/_{2}
(D) 6 0 1 ^{1}/_{2}
Emission Spectra Lab
PreLab Questions:
1. According to Bohr's atomic model, where may an atom's electrons be found?
2. How do electrons become excited?
3. State the equation that is used to determine the energy content of a packet of light of specific frequency.
4. What form of energy emission accompanies the return on excited electrons to the ground state?
Write and/or draw your observations as you view the emission spectra.
DATA TABLE
Gas 
Observations 
incandescent 

hydrogen (H_{2}) 

carbon dioxide (CO_{2}) 

helium (He) 

neon (Ne) 

water vapor (H_{2}O) 

air 

mercury (Hg) 

argon 

krypton 

xenon 

nitrogen
(N_{2}) 

iodine
(I_{2}) 

oxygen
(O_{2}) 

fluorescent 

FLAME TESTS FOR METALS LAB
Background: The active metals of groups 1 and 2 can be "excited" in a flame. The energy (in the form of heat) in the flame causes the electrons in the metal to jump up into higher energy levels. When the electrons fall from the excited state, they produce light. Each metal produces a characteristic color of light.
Purpose: To identify the presence of a metal found in each solution by observing the color produced when metal compounds are excited in a flame. To determine the identity of a metal ion in an unknown solution.
Lab
Safety:
**
ALWAYS WEAR YOUR SAFETY GOGGLES! **
** TIE BACK LONG HAIR! **
Procedure:
1.
Select one wooden splint from the container for the element you are testing.
2. Place it into the flame as demonstrated by your instructor. Place burned wooden splints into beaker of water.
3. Carefully observe the color of the flame and record your observations.
4. Test the remainder of the solutions.
5. Compare the known solutions with the unknown solution and record your observations.
6. Clean up your lab station as directed by your instructor.
Data: Record the color of the flame for each of the known solutions.
Metal Ion 
Color of Flame 
Lithium 

Sodium 

Potassium 

Calcium 

Barium 

Copper 





Unknown # ____ 

Questions:
1.
What is the identity of the unknown based on your observations? How did
you know?
2. According
to the Bohr model of the atom, what happens in the atom that causes colors to
be emitted during these
flame tests?
3. Why should you use a separate wooden splint for each element you test? (Why not reuse a partially burned wooden splint?)
4. What do you think would happen if the unknown substance contained a mixture of two compounds? Could each metal be identified?
5. Understanding the flame test properties of the
group 1 and 2 metals, what application could these chemicals be used for?