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Honors Chemistry – Unit 3 Bohr’s Model, Atomic Structure, Electron Configuration, Nuclear Chemistry
Scientist Quiz: Friday, Feb. 20
Atomic History Timeline Project: Friday, Feb. 27
Test Date: Tuesday, March 3
Vocabulary:
isotope
wavelength
mass number
electromagnetic radiation
proton
neutron
orbitals
excited state
Aufbau Principle
valence electrons
Law of Definite Proportions
Law of Multiple Proportions
electromagnetic spectrum
Vocab Quiz: Tuesday, Feb. 24
EQs/PS Due: Tuesday, March 3
average atomic mass
electromagnetic spectrum
electron
ground state
Pauli Exclusion Principle
Law of Conservation of Energy/Mass
Heisenberg’s Uncertainty Principle
Important Scientist Who Contributed to Atomic Theory:
Democritus
Dalton
Thomson
Moseley
Pauli
Heisenberg
Schrodinger
Millikan
Planck
Bohr
Chadwick
Formulas/constants will be given on test:
Average atomic mass problems
atomic number
frequency
ion
quantum numbers
Hund’s Rule
DeBroglie
Hund
Rutherford
Radioactive Decay problems
Bohr Model
OBJECTIVES:
 Given a scientist’s name be able to describe his or her contributions to chemistry and the development of atomic structure and/or
the periodic table.
 Be able to find the number of protons, electrons and neutrons for different isotopes.
 Understand how to correctly write an isotope.
 Be able to differentiate between alpha (α), beta (β), and gamma (γ) radiation and decay.
 Be able to draw and describe the basic structure of an atom.
 Be able to do average atomic mass problems
 Be able to find wavelength (using Bohr’s model of H)
 Understand the relationship between wavelength, frequency and energy
 Be able to list and describe the four quantum numbers.
 Be able to write electron configurations using the periodic table.
 Be able to explain the uncertainty principle, Hund’s Rule and Pauli’s exclusion principle
 Be able to write Noble gas configurations, orbital diagrams and dot diagrams
 Understand the wave-particle duality of nature
 Be able to define and correctly use vocabulary above.
 Be able to explain the law of conservation
 Understand the relationship between wavelength, frequency and energy
EQS Unit 3:
Day 1 - Review:
 Calculate the density of a 22.0 g of a metal with a volume of 4.9 cm3? Identify the metal using your reference pack.
 Convert 250 ml to L.
Day 2 - New Material:
 Draw an atom – label its parts (protons, neutrons and electrons)
 Describe the relative mass and relative charge on each subatomic particle.
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Day 3 –
 Describe the contributions that the following individuals made to the atomic theory:
a. Dalton
d. Bohr
b. Thomson
e. Rutherford
c. Moseley
f. Schrodinger
 Copy and fill in the following table:
Element/ion
Fe
K+
# of protons
27
# of neutrons
# of electrons
25
O2Strontium-89
Day 4  Find the mass of an element if out of a sample of 100:
o 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?
o Identify this element by symbol and name.
Day 5 –
 Write the long form (full) electron configuration for arsenic.
 Write the Noble Gas electron configuration for
o Al
o Ag
o At
Day 6 –
 Give the Lewis dot diagram for the following:
o P
o Ba
Day 7  Explain how a glow in the dark sticker works – use these words in your explanation- ground state, excited, electrons,
jump(ing).
 Using Bohr’s model:
o what is the wavelength of light when an electron jumps from n= 6 to n = 3?
o What type of energy is this?
Day 8  Comparing Blue light, X rays and Radio waves which has
a. the larger wavelength?
b. the highest frequency?
c. the most energy?
Day 9 –
 Complete the following nuclear reactions:
a.
226
88
Ra  ?? ? +
b.
209
84
Po 
205
82
0
1
e
Pb + ?? ?
Day 10 –
 Which of the three radioactive emissions ( best fit the following statements? Write the correct symbol/s on the lines.
Some may be used more than once.
a. These emissions are charged.
b.
This emission is the most massive (heaviest).
c.
This emission is the most charged.
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Atomic theory How was the atomic theory developed?
_____________________________: first to use the term atom
Dalton: (1808) Atomic Theory: elements are composed of extremely small particles called ______________

Atoms of the same element are __________________ in size, mass and properties

Atoms cannot be __________________ or _______________________ (Law of Conservation of Mass)

Atoms of different elements ____________________________________________________ to form compounds

In _________________________ reactions, atoms are combined, separated or rearranged.
JJ Thomson (1897) Discovers the _________________________

Proposes the “_______________________________” model of the atom:

Atom is neutral and electrons are embedded in a sphere of positive charge.
Max Planck (~ 1900) Developed the theory of _____________________________
_____________________________ (~ 1905) light behaves as a particle called a ______________________
Milikan (1909) discovers electron charge and mass. (through his _________________________________).
Rutherford (1911) {was a student of Thomson’s; believed in the plum pudding model)___________________________ (a thin piece
of gold foil was bombarded with a narrow beam of alpha particles, the majority passed through undisturbed, but a small amount were
deflected by the nucleus). He discovered that the

Atom is mostly _________________________

Atom has an extremely small, dense and ______________nucleus.

Electrons move around the nucleus. (Not embedded in the nucleus)
*Bohr (1913) Electrons circle nucleus in specific circular paths at fixed distances from the nucleus. Each electron orbit has a specific
energy or _______________________. A quantum of energy is required to move from one orbit to the next. Works for hydrogen
atom only. We will do work with Bohr’s model in the next section!
DeBroglie (1923) ____________________________________________ - proposes wave-particle behavior of electrons
Werner Heisenberg (~1927) - ___________________________________ - states that we cannot know both the position and the
momentum of a quantum particle (electron) at any given time
Schrodinger (~1928) writes an equation to determine the probability of electron location called the __________________
_____________________________
Chadwick (1932) discovers the neutron
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*Electron Cloud model (present) based on Schrodinger’s wave equation.
 Visual model of probable locations of the electrons in an atom. Atoms are composed of electrons in a cloud around a
positive nucleus.
Laws:
1. Law of conservation of mass (matter):
2. Law of definite proportions: a chemical compound contains the same elements in exactly the same proportions, by mass,
regardless of size or source of sample.
3. Law of Multiple Proportions: If 2 or more different compounds are composed of the _______________________, then the
ratio of the 2nd. element combined with the a certain mass of the 1 st. element is always a _______________ of small whole
numbers
Weighing and Counting Atoms
Subatomic particles:
Particle
Location
Relative charge
Relative mass
Proton
Neutron
Electron
Atomic # (symbolized with Z) = # of protons
Identifies (ID’s) an element – UNIQUE for each element
Elements in order on the periodic table by atomic #
Examples: H  1;
He  2, etc.
Ag = ______
28 = _______
Because atoms must be neutral; atomic number also = # eTherefore Na has ______ p+ and _______ eMass # (often symbolized with A) - total # of p+ and no (electrons are not include they have an insignificant mass)
Therefore # no = mass # - atomic number
(rounded to nearest whole #)
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How many p+, no, and e- in F? p+ = ______ , no =__________, and e- =_________
Example:
You try:How many p+, no, and e- in K? p+ = ______ , no =__________, and e- =_________
Ions = Atoms of the same element with a different number of _____________ which leads to the element having a
______________
charge occurs from gaining or losing e- (not p+)
Cation: + = lost eAnion
− = gained eBa2+
example:
you try Cl-1
p+ = ______ ,
no =__________, and e- =_________
p+ = ______ , no =__________, and e- =_________
Isotope = Atoms of the same element with a different number of _______________ which leads to the element having a
different __________
Isotopes can be written in two forms:
mass #

Uranium-235
235
92 U

element
or

mass #
 element symbol

atomic #
Example for Uranium-235 p+ = ______ , no =__________, and e- =_________
You try: for Nitrogen -15 p+ = ______ ,
no =__________, and e- =_________
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Atomic Structure Summary sheet
Most people already know that the atom is made up of three main parts, the _______________ and ______________ in the nucleus
and the ______________ somewhere outside of the nucleus.
How Many Particles in Each Atom?
The particle that defines the identity of an atom is the _____________. (shown on the periodic table)
Every hydrogen atom has ___ proton.
Every magnesium atom has ___ protons.
Any atom that has 23 protons is _________________.
Any atom that has 92 protons is _________________.
The mass of an atom is mostly from the ___________ and ____________.
Find O on the periodic table. It’s mass is ______ amu.
It has ___ protons. It must have ___ neutrons.
If the mass is not close to a whole number, it is because the atom has several _____________. These are atoms with the same number
of ___________ but different numbers of _____________.
Chlorine has two isotopes: Cl-35 ( ___ p+ & ___ n) and Cl-37 ( ___ p+ & ___ n).
If an element has a charge it is called an ____________ and the charge is due to the loss or gain of ______________. For example the
sulfur atom has _____ p+, ________ n0 and ________ eBut the sulfide ion: S-2 has _____ p+, ________ n0 and ________ e-
Fill in the following table:
Atom
He
Si
Sr2+
H
Rn
Argon-39
F-1
lead-208
protons
neutrons
electrons
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Average Atomic Mass
Atomic mass: relative scale based on Carbon-12
Unit for atomic mass = atomic mass unit = amu or u
Why are there decimals on the masses on the periodic table?
Calculating average atomic mass:
(mass) x (%) + (mass) x (%) + (mass) x (%) etc
100
Example: What is the average atomic mass for carbon if 98% is Carbon-12 and 2 % is Carbon –14?
You Try: What is the average atomic mass for Bromine if 50.69 % has a mass of 79 amu and 49.31 % has a mass of 81 amu?
What’s the Difference between Mass Number and Average Atomic Mass?
http://www.youtube.com/watch?v=m15DWkkGe_0
What is the mass number?
How much do a proton and a neutron weigh?
Why do we not concern ourselves with the mass of electrons when determining the weight of an atom?
How do we name the isotope of Boron that has 5 neutrons?
Atomic Mass: Introduction
http://www.youtube.com/watch?v=7fYpEnxhKQk
How do you change a percent into a decimal?
What does a weighted average take into account that a regular average doesn’t?
How do you calculate the weighted average of an element?
Why didn’t Mr. DeWitt’s first weighted average of copper match the average atomic mass on the periodic table?
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Electron Configuration
Electrons have properties of both waves and particles = wave-particle duality
Remember the Bohr model for hydrogen:
- An electron circles the nucleus in orbits – (fixed energy ranges)
- An electron can gain energy- become excited - and go up an orbit (or more) or it can emit energy and move down an orbit (or
more).
- The lowest energy orbit (the ground state) is closest to the nucleus
Additional info relating to electrons:
A quantum is a packet of energy (this is the energy required to move an electron)
Heisenberg uncertainty principle states that it is impossible to know both the position and velocity of an electron at the same time.
An orbital is the 3-d region around the nucleus that indicates the probable location of an electron.
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Quantum numbers describe electrons and their orbitals – we will study quantum numbers next.
QUANTUM NUMBERS AND ATOMIC ORBITALS
Quantum # (Q#) = #s that specify the properties of atomic orbitals and their electrons.
There are 4 Quantum #s
1. Principal Quantum Number: (n)
 Indicates energy levels (shells), n = 1, 2, 3 etc. Diagram =
 As n increase so does energy
 # e-s on a level = 2n2 Example: e-s on energy level 2?
You try: How many electrons on energy level 4?
2. Orbital Quantum Number: (
)

indicates the shape of an orbital and the subshells or sublevels

subshells: s, p, d, f

n = 1 , energy level one has 1 subshell named 1 s

n = 2, has 2 subshells named 2s, 2p

n= 3 has ________
subshells named __________

n= 4 has _______
subshells named _____________
each subshells contains orbitals (think of them as the electron’s home)
s has 1 orbital, p has 3 orbitals, d has 5 orbitals, f has 7 orbitals.
Each orbital can hold a maximum of 2 e-s.
Therefore:
s = 1 orb,
2
e-
p = ______
orbs,
_____e-
d = _____
orbs,
______ e-
f = ______
orbs,
_______e-
Shapes of orbitals:
3. Magnetic Quantum Number ( ): indicates orbital orientation about the nucleus.
See example on board:
4. Spin Quantum number: ( ):
indicates spin of e2 values = +1/2 and – ½ (clockwise or counterclockwise)
Pauli’s exclusion principle: no two electrons in the same atom can have the same 4 quantum numbers – therefore electrons in the
same orbital must have opposite spins.
Energy Level
n=1
n=2
n=3
n =4
Subshells
# of Orbitals
Maximum number of
e -s
Total number of e-s
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Electron Configuration – the arrangement of electrons in atoms
 Remember, atoms of different elements have different number of electrons.
 Use the atomic number to determine the number of electrons for an atom.
For example:
Na has ____e-;
S has _____e-
The Aufbau Principle – tells us an e- occupies the lowest energy orbital that can receive it.
To fill orbitals: Begin first filling the 1s (2e-s), then the 2s (2e-s), 2p (6e-s), then 3s (2e-s), and 3p (6e-s).
Electron configurations tell us:
3p5
Energy level
Examples:
exponent = # of e-s
Letter = sublevel
H
He
Li
S
You try: nitrogen
Silicon
When we get to the 3d and 4s orbitals, we see some overlapping of orbitals.

This is due to the overlapping of energy levels

This overlapping causes the 4s orbital to fill before the 3d orbital.

There is even more overlapping in the 4th & 5th , 5th & 6th, etc. sublevels.

Therefore, the 4s orbital will “fill up” with e- before 3d.
Shortcut for filling orbitals: Use periodic table – we will label together.
Using the periodic table write the following electron configurations
K
You try: P
Ex: Br
You try: Co
Example: Pt
You try: Pb
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Before we talk about other type of notations we have to discuss:
valence electrons (outer electrons) – electrons in the outermost energy level. (participate in bonding!)
inner-shell electrons – electrons not in the highest occupied energy level.
Examples: Circle the valence electrons in the electron configurations on previous page.
1. Noble Gas Notation: a shorthand way to write electron configuration.
noble gases – group 18 – have full outer energy levels (8e-s or 2e-s for He).
We can use this to shorten electron configurations:
ex.
Na
Ex. 2:
Ne
Ti
You try:
Cl
Y
and
2. Orbital Notation: an orbital is represented by
W
and electrons are represented by
or
(arrows).
Hund’s Rule: Empty orbitals fill before electrons begin pairing up:
NOT
2p3
2p3
You try the orbital notation for:

Se

Ni

Na
3. Electron Dot Diagram: a third way to represent electron configurations using valence electrons.
Step 1: Draw e- configuration & orbital notation.
Step 2: Allow element symbol to represent inner e-s.
Step 3: Use dots for valence e-s.
a)
It is important to pair e-s correctly.
b)
It is not important which side of the element symbol the dots are placed.
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H
Se
Pd
You try:
Mg
Sb
Xe
Electron configuration practice sheet
Write the long form electron configuration for the following elements:
1.
Sodium ____________________________________
2. Iron
__________________________
3.
Bromine ___________________________________
4. Barium
_________________________
5.
Gallium ___________________________________
Write the noble gas electron configuration for the following elements:
6. Cobalt
________________________
7. Silver
___________________________
8. Tellurium
_______________________
9. Radium
__________________________
Determine what elements are denoted by the following configurations
10. 1s22s22p63s23p4
________________
11. [Kr] 5s24d105p3
________________
12. 1s22s22p63s23p64s23d104p65s1 ________________
13. [Xe] 6s24f145d6
________________
14. [Rn] 7s25f146d4
________________
Determine which of the following configurations are not valid; explain why not.
15. 1s22s22p63s23p64s24d104p5 ________________________
16. 1s22s22p63s23d5 ________________________
17. [Ra] 7s1_______________________________________
18. [Kr] 5s24d105p5 ________________________
19. [Xe] _________________________________________
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Radioactivity Ppt Notesheet
 The presence of too many or too few ____________________________ relative to protons can cause a nucleus to be
________________________________.
 An unstable nucleus will _______________________ to become stable.
 This process of _______________________________________ releases a lot of energy.
 Also called __________________________________________________________.
 Radioactivity produces both penetrating __________________ and ____________________.
 The particles and rays are called __________________________________________.
 3 Types of Radiation: ____________________, _____________________, ____________________
Alpha Radiation
 ___________________ nuclei that have been emitted from a radioactive source.
 These particles are called ________________ particles.
 Contain ________ protons and _________ neutrons—very ____________________.
 Have a ______________ positive charge. ++
Beta Radiation
 An ___________________ resulting from the breaking apart of a neutron in an atom is called a ________________ particle.
Gamma Radiation
 A high-energy photon emitted by a radioisotope is called a gamma _____________.
 Often emitted ___________ either alpha or beta particles.
 Has no ______________ or __________________.
Penetrating Power
 Alpha particles are the _______________ penetrating. (They are stopped by ___________________.)
 Beta particles are stopped by _______________ or aluminum foil.
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 Gamma particles are the _______________ penetrating. (They are not even completely stopped by ____________ or concrete.)
Practice:
1. An unstable nucleus releases energy by _________________________________________________.
2. Which property does NOT describe an alpha particle?
a. 2+ charge
b. a relatively large mass
c. a negative charge
d. low penetrating power
3. When a radioactive nucleus releases a high-speed electron, the process can be described as
____________________________________________.
Half-Life
 A half-life (t1/2) is the time required for one-half of a radioisotope sample to ______________ to products.
 After each half-life, half of the existing radioactive atoms have decayed into atoms of a ___________ element.
Show your work.
1. Carbon-14 emits beta radiation and decays with a half-life (t1/2) of 5730 years. Assume you start with a mass of 2.00 x 10-12 g of
carbon-14.
a.
How long is three half-lives?
b.
How many grams of the isotope remain at the end of three half-lives?
2.
Manganese-56 is a beta emitter with a half-life of 2.6 h. What is the mass of manganese-56 in a 1.0 mg sample of the isotope at
the end of 10.4 h?
3.
If there were 128 grams of radioactive material initially what mass remains after four half-lives?
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Radioactivity Worksheet
1. How do you find the amount of radioisotope remaining after decay if the number of half-lives is given?
2.
How do you find the amount of radioisotope remaining after decay if the number of half-lives is not given?
3.
A patient is administered 20 mg of iodine-131. How much of this isotope will remain in the body after 40 days if the half-life
for iodine-131 is 8 days?
4.
State the number of neutrons and protons in each of the following nuclei:
5.
a.
2
1
b.
56
26
c.
197
79
H : ________________________________________________________
Fe : _______________________________________________________
Au : ______________________________________________________
The three types of radioactive emissions are called alpha (), beta () and gamma () radiation. Complete the table below
with the correct information about each type.
Charge
Atomic Symbol
Can Be Stopped By
Alpha
Beta
Gamma
6.
7.
8.
n/a
Which of the three radioactive emissions ( best fit the following statements? Write the correct symbol/s on the lines.
Some may be used more than once.
a.
These emissions are charged. ____________
b.
This emission is the most massive (heaviest). ____________
c.
This emission is the most charged. ____________
Which type of radiation – alpha, beta, or gamma:
a.
Results in the greatest change in atomic number? Why?
b.
Produces the greatest change in mass number? Why?
Complete the following nuclear reactions:
a.
238
92
U  ?? ? + 24 He
b.
234
90
Th 
234
91
Pa + ?? ?
________________________________________________________
________________________________________________________
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9.
218
84
When
Po emits a beta particle, it transforms into a new element.
a. Write out the nuclear equation:
b.
Fill out the chart below:
Name of the
Element
Parent
Element
Daughter
Element
Atomic
Number
Mass #
# Of
Protons
# Of
Neutrons
Polonium, Po
Astatine, At
10. How does the radioactive isotope C-14 differ from its stable counterpart C-12?
11. Determine the average atomic mass if chlorine exists in the following mixture of isotopes:
75.77%
35
17
Cl
and
24.23%
37
17
Cl
Electromagnetic radiation:
Light as a wave form. Any wave is essentially just a way of shifting energy from one place to another - whether the transfer
of energy in waves on the sea or in the much more difficult-to-imagine waves in light. In waves on water, the energy is
transferred by the movement of water molecules. But a particular water molecule doesn't travel all the way across the Atlantic
- or even all the way across a pond. Depending on the depth of the water, water molecules follow a roughly circular path. As
they move up to the top of the circle, the wave builds to a crest; as they move down again, you get a trough.
Wavelength, frequency and the speed of light. If you draw a beam of light in the form of a wave (without worrying too
much about what exactly is causing the wave!), the distance between two crests is called the wavelength of the light. (It could
equally well be the distance between two troughs or any other two identical
positions on the wave.) You have to picture these wave crests as moving from left to right. If you counted the number of
crests passing a particular point per second, you have the frequency of the light. It is measured in "cycles per second”, which
is called Hertz, Hz. Orange light, for example, has a frequency of about 5 x 10 14 Hz . That means that 5 x 1014 wave peaks
pass a given point every second. Light has a constant speed through a given substance. For example, it always travels at a
speed of approximately 3 x 108 meters per second in a vacuum. This is actually the speed that all electromagnetic radiation
travels - not just visible light.
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The relationship between frequency and wavelength can be expressed with the equation: C =  C = constant = speed of
light 3.0 x 108 m/s  = (lambda) wavelength in m or nm  = (nu) = frequency in /sec or Hertz (Hz). This relationship
means that if you increase the frequency, you must decrease the wavelength.
Compare this diagram with the similar one above. If the wavelength is longer, the frequency is lower.
Consider the diagram below (Consider it wave One) in order to answer questions.
1. The wavelength of the wave in the diagram above is given by letter ______
2. Below draw a similar wave with a smaller wavelength than the one above (this will be wave Two)
Is wave two’s frequency higher or lower than wave one’s frequency? _____________
3. Below draw a similar wave with a bigger wavelength than wave one (this will be wave Three)
Is wave three’s frequency higher or lower than wave one’s frequency? _____________
4. Considering all three waves: which has: A. The highest frequency?____________
B. The lowest frequency? _______________
C. The largest wavelength?_______________
D. The smallest wavelength ?______________
From the above observations we can conclude that a wave with a large wavelength will
have a _______________ frequency and a wave with a small wavelength will have a
________________________ frequency.
Additional info: The study of substances exposed to energy is spectroscopy. The energy released by these substances are
spectrums. Visible light (ROY G BIV) is only part of the spectrum. The whole spectrum is shown on page 92 (or pg139).
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Planck proposed that a _quanta_ is a packet of energy emitted in small, specific amounts- when the quanta is light it is
often referred to as a photon. Einstein expanded on Planck’s idea by introducing the idea that radiation has a dual-wave
particle nature(often referred to as wave-particle duality) ie light (and electrons) has properties of both waves and
particles.
Bohr used the quantum theory to describe hydrogen: an e- circles the nucleus – the e- can absorb energy and jump up
a level (or more). When the e- drops back down it emits the extra energy in a quanta (as light or heat).
The lowest energy state an electron can occupy, the one closest to the nucleus is called its ground state (think of it as the
electron’s home).
When atoms are exposed to energy they will emit certain colors (certain frequencies of visible light). Different elements have
electrons that will jump different amounts of orbits, producing different colors. For example gas tubes in neon signs are
special elements. Pink tubes contain neon and blue tubes contain argon. Spectroscopes are used to study different emissions.
Bohr’s Model of H
Refer back to Bohr’s model in reference packet!
Electrons jump energy levels.
Up = gaining energy - this is referred to as the excited state
Down = releasing energy – this is referred to as the ground state (home- lowest energy) energy is released in the form of
light.
Equation: We will not work problems but you need to understand that wavelength and frequency are inversely related if
one increases the other decreases.
C = 
C = constant = speed of light 3.0 x 108 m/s
 = (lambda) wavelength in m or nm
 = (nu) = frequency in /sec or Hertz (Hz)
E = h We will not use this equation in here – but it will be used in higher level classes. You do need to understand that
Energy and frequency are directly related; higher freq = higher Energy!
E = energy in Joules (J)
h= Planck’s constant = 6.63 x 10 –34 J-sec (given)
 = (nu) = frequency in /sec or Hertz (Hz)
Example problems: Use Bohr’s model
1. What is the wavelength of light absorbed when a hydrogen electron jumps from n = 1 to n =3? What type of energy
is this?
You try: 1. What is the wavelength of light absorbed when a hydrogen electron jumps from n = 3 to n = 4? What type of
energy is this?
Example: 2.What color of light is released when a hydrogen electron jumps from n = 4 to n =2?
You try: What color of light is released when a hydrogen electron jumps from n = 3 to n =2?
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Understanding the model (together):
Between gamma, X-rays, and blue light
1. Which has the longest wavelength? ____________
2.
the shortest? ______________
3.
Which one has the highest frequency? _______________
4.
The lowest energy? _______________________
Practice
1. What wavelength of light is released when a H e- jumps from n=4 to n=1
2. What type of energy is this?
3. What is the wavelength of light released when a H e- jumps from n = 4 to n = 2?
What type of energy is this?
4. Comparing UV, red light and IR – which one has the shortest wavelength?________
The highest frequency?________________
highest energy? _______________
5. What color of light is released when an e- jumps from n = 6 to n = 2?
Bohr Model Worksheet Extra practice
1. In each pair below, which form of radiation has more energy?
a. Red light or Violet light?
c. Radio waves or Gamma rays?
b. Infrared waves or Ultra-violet rays?
2. Which of the following actions would cause light to be emitted?
a. an electron jumps from the n=2 to the n=4 energy level.
b. a natural gas source is ignited and burns.
c. an atom absorbs energy from heat.
d. an electron falls from the n=3 energy level back to the ground state.
3. What is the wavelength of energy produced when an electron falls from n = 2 to n = 1; what type of energy is this?
4. Black lights make use of long-wave ultraviolet light, usually with a wavelength around 365 nm (they also give off some
purple-colored visible light) and is generally harmless. Short-wave ultraviolet light around 254 nm is used in the goggle
sanitizer units in our labs, and would damage our eyes if we looked at it directly. Which of these light types has a higher
frequency? Which is higher in energy? How does their energy correspond to their respective uses?
Atomic theory packet - Page 21 of 28
MINI LAB – Average Atomic Mass of “Candyium”
Please do NOT open the bag until you receive permission. Working in groups of three to four, complete the following
activity.
EVERY student in the group must do the work for this problem in their notebook. No group will receive credit for this
assignment unless every member of the group works the candyium problem! Each group will receive a baggie labeled A, B,
C, or D. Do not open the bag, until you are given permission to do so. Inside the bag are ten “atoms” of the new element
“candyium”. Your mission is to calculate the average atomic mass of your bag of candyium using the data below.
If you are successful in your mission – your reward will be the opportunity to more closely examine and keep the candyium
(ie. eat it!). After everyone in your group completes their calculations, raise your hand and allow Mrs. Reid to check your
math (include units); if your work is correct, you may have the candyium.
DATA:
Color of isotope
Green
Red
Yellow
Orange
Purple
Mass (u)
7
6
5
8
9
Each person should record the following on their own packet:
Data Table
Bag letter:____________
Isotope color
number in bag (out of 10)
percentage
Green
__________
_______
Orange
_________
________
Red
_________
________
Purple
_________
_________
Yellow
__________
_________
Work for calculation of average mass for your bag: Show Work below
Average mass (with units):____________________
Mini-lab Exploring the concept of the law of conservation of mass
Copy the bold parts of the lab onto your own paper!
Objective: to explore changes in mass that may occur during chemical or physical changes.
Pre-lab Questions: 1. What is the law of conservation of mass (matter)? ______________________
_____________________________________________________________
Procedure:
1. Put a scoop of baking soda (NaHCO3) in one cup. Put 10 ml of 1.0 M dilute acetic acid/vinegar (HC2H3O2) in the
other cup.
2. Place both cups in a sealed Ziploc bag and place the bag on the balance; record the total mass.
3. Combine the reactants (take them off the balance to combine them – pour the liquid into the cup with the baking
soda) and record observations.
Atomic theory packet - Page 22 of 28
4.
5.
6.
Record the total mass after the reaction – put the whole bag back on the balance.
Dump the mixture down the drain – rinse your cups with water and continue answering questions.
Leave out supplies for other classes – unless you are the last class; the last class will follow clean up directions! (see
board)
Data table: COPY THIS TABLE ONYOUR OWN PAPER!
Mass of reactants (before
mixing)
Mass of products (after
mixing)
Observations
Questions: answer on your paper!
1.
Was there a difference in the mass of the reactants and the mass of the products?
2.
How does this lab demonstrate the law of conservation of mass?
3.
What sign of a chemical change did you observe?
4.
Predict what your product(s) are. Make a prediction!!
Honors Atomic Structure Review Sheet
A. Completion:
Use this completion exercise to check your knowledge of the terms and your understanding of the concepts introduced in this unit.
Each blank can be completed with a term, short phrase, or number.
Atoms of each element are (1) ______________________ from the atoms of all other elements. Dalton theorized that atoms are
indivisible, but the discovery of subatomic particles changed this theory. We now know that atoms are made up of electrons, which
have a (2) _________________ charge; (3)___________________ , which have a positive charge, and (4)_____________________ ,
which are neutral. The latter two particles are found in the (5) _______________ of the atom.
It was (6) __________________________who discovered the nucleus of the atom. The nucleus has a positive charge and it occupies a
very small (7)______________________ of the atom. In contrast, the negatively charged (8)_______________________ occupy most
of the volume of the atom. The number of (9)_______________________ in the nucleus of the atom gives the atomic
(10)_______________________ of that element. Because atoms are electrically neutral, the number of protons and (11)
____________________are equal. The sum of the
(12) ______________________and neutrons is the mass number. Atoms of the same element are identical in most respects, but they
can differ in the number of (13)____________________________ in the nucleus. Atoms that have the same number of protons but
different mass numbers are called (14)________________________.
(15) ________________________ are atoms that have the same number of protons but a different number of electrons resulting in a
(16)___________________________. The (17)_________________________ of an element is the weighted average of the masses of
Atomic theory packet - Page 23 of 28
their isotopes of that element. An atom of sulfur-32 isotope contains (18) ____________________________ protons and (19)
___________________________neutrons.
In the Bohr model, the (20) ____________________________ move in paths of fixed orbits. (21) ______________________ is the
distance between successive peaks of a wave and is measured in (22)___________________________.
(23) _____________________is the number of waves per second and is measured in (24)________________________.
Wavelength and frequency are (25) ____________________________related; meaning if a wave has a large wavelength it has a
(26)_____________________________ frequency.
1. Given the relative abundance of the following naturally occurring isotopes of bromine, calculate the average atomic mass.
Br-79 53%
Br-81 47%
2. What is the color of light that has a wavelength of 4.7 x 10 -7 m?
3. Fill in the following table:
Element/ion
Atomic #
Atomic
Mass
Mass
Number
Protons
Al
S-2
33
16 S
Ba2+
4. What is the wavelength, color and energy type of light that jumps from n = 6 to n = 2
5. Which has the most energy: microwaves, blue light or x-rays?
6. How is light produced (include electrons in your explanation!
7. Fill in the following table:
Element
F
Cl
Ca
K
Cu
O
Ne
metal or nonmetal
period
family
Neutrons
Electrons
Atomic theory packet - Page 24 of 28
FLAME TEST LAB
When elements are heated to high temperatures, their atoms are placed in an excited state. In the excited state, electrons have absorbed
the exact amount of energy needed to “jump” outward from their ground state to a higher energy level. The exact energy needed is
determined by the distance between the energy levels and is unique to each element. The excited electrons then release the absorbed
energy and return to a lower energy level. We are able to see that emitted energy if it is in the visible spectrum. The observed colors,
or “spectra”, the substance emits is a unique set of wavelengths from low to high frequency (ROYGBIV), and the emission spectra
can therefore be used to identify the element.
In this experiment you will perform a flame test on several metallic salts that are dissolved in water. As the metal ions are heated you
will be able to see the dominant color in the element’s spectrum. Based on your observations, you will develop a reference table that
lists the flame color for each metal ion. You will then perform flame tests on four unknown metallic salts and by comparison to the
reference chart be able to identify them.
Procedure
1.
Follow the proper procedure to light your burner. Once lit, one partner is to remain on watch at all times.
2.
Dip an inoculation loop into a labeled solution (not an unknown) located in your test tube rack. Hold it upright in the hottest part
of the flame, and record your observations in the data table located on the back of this lab sheet.
3.
Rinse loop in a weak acid and then in distilled water. Repeat step 2 for each of the known solutions and then for the unknown
solutions.
4.
If you are uncertain of the identity of an unknown, retest the solutions to clarify.
5.
Clean lab area carefully!
Solution
Flame color
Solution
Sodium chloride
Strontium
chloride
Barium chloride
Lithium chloride
Calcium chloride
Copper chloride
Flame color
Potassium
chloride
Color
Unknown A
Unknown B
Other observations
Identity
Atomic theory packet - Page 25 of 28
Analysis Discussion Questions: (Analysis Discussion Questions must be answered in complete sentences and provide details to support
statements. This portion counts for 50% of your lab grade. This will be included in the Analysis portion of your lab report.)
1.
Which elements are easiest to identify? What makes them easy? (Make sure to discuss more than one element.)
2.
Which elements are most difficult to identify? What makes them difficult? (Discuss more than one element.)
3.
4.
Do you think flame tests would be reliable for detecting metal ions in a mixture? Explain why or why not.
Explain any inconsistencies or errors in your lab. Provide explanations for these errors.
Atomic theory packet - Page 26 of 28
Problem Set Unit 3
You must show work for credit – put final answer on the blank provided include units!
Classify the following as being a physical or chemical change.
Write P or C on blank!
1. Sugar dissolves in water.
________
2. Hydrochloric acid reacts with potassium hydroxide to produce a salt, water and heat.
________
3. Water is heated and changed to steam.
_________
4. Iron rusts
________
5. Ice melting
________
6. Wood rotting
________
7. Water is absorbed by a paper towel.
________
Classify the following as a solution, element, Heterogeneous mixture (HM) or compound.
8. Chlorine gas
_______________
10. rocky road ice cream _______________
9. Steam _______________
11. Sugar water _______________
Math problems Show work!
12. You are driving 55 miles per hour how many meters will you travel in 20 minutes?
_______________________
13. A piece of metal that is 2.0 cm long, 3.2 cm wide, and 10.0 cm high is melted.
a. What is the liquid volume in mL?
a. ________________
b.
in L?
b. ________________
14. In March 1989, the Exxon Valdez ran aground and spilled 240,000 barrels of crude petroleum off the coast of Alaska. One barrel
of petroleum is equal to 42 gallons. How many liters of petroleum were spilled?
________________
Atomic theory packet - Page 27 of 28
15. The density of dry air measured is 0.00119 g/cm3. What is the volume of 400.0 g of air?
_______________
16. A 25.0 g sample of a metal has a volume of 2.21 cm3 – Find the density of the metal and identify the metal.
D = ____________________
ID = ____________________
17. Given microwaves, yellow light and gamma rays which has
a. the shortest wavelength
___________________
b. the lowest frequency
___________________
18. Draw and label an atom (include all the parts with charges) Provide a key to identify your subatomic particles.
19. Calculate the average atomic mass for copper if 69% has a mass of 63 amu and 31% has a mass of 65 amu.
____________________
20. Find the wavelength of light if an electron jumps from n = 5 to n = 3.
____________________
21. Find the number of protons, neutrons, and electrons in the following:
p+
no
ea. Sodium
____
_____
______
b.
89
38Sr
_____
_____
______
c.
Aluminum-25
_____
_____
______
d.
Ba2+
_____
______
______
e.
Iron-57
______
______
_____
Atomic theory packet - Page 28 of 28
22. Radioactive iodine-131 has a half-life of eight days. The amount of a 200.0 gram sample left after 32 days would be — (Show all
work)
23. When isotope bismuth-213 emits an alpha particle:
a.
Write out the nuclear equation:
b.
What new element results if the isotope, instead, emits a beta particle? (Write out the equation here.)