Download Assignment # 6 Atomic Structure Drill

Chemistry 112A
Mrs. Ozment, Mr. Harvey & Mrs. McCall
UNIT 4:
Atomic Concepts Structure
Student Study Guide
Name:
1
UNIT 4: Atomic Structure
Unit 4 Study Guide: Atomic Concepts & Periodic Table
Essential Questions for the Unit:
 How has the view of the atom evolved over time?
 What are the parts of an atom, where are they located, and how are the
parts held together?
 What are isotopes?
 What is the atomic number for an element?
 What is an ion?
 How is an amu calculated?
 What are electron energy levels and how are they used in chemistry?
Vocabulary:
atom
nucleus
subatomic particle(s)
 Proton
 Electron
 Neutron
gold foil experiment
electron energy levels
atomic number
atomic mass
atomic mass unit (amu)
ion
isotope
elements
chemical property
chemical change
What you need to know and be able to do within Unit 4 is listed at the top of each topic page
**TOPIC 1**
Development of the Atomic Model
You Should Know/Understand OR Be Able To Do:
1. Define an atom as the smallest piece of matter retaining elemental properties.
2. Summarize Dalton’s Atomic Theory and relate this theory to models of atoms and compounds learned
in Unit 1.
3. Describe the atomic theories of Thomson, Rutherford, and Bohr and relate experimental evidence to the
model of the atom.
4. Discuss the structure of an atom including location of the protons, electrons, and neutrons in relation to
the nucleus
2
NOTES:
Atom: smallest particle of an element that retains its properties.
Atoms are so small we have special microscopes to see them.
These pictures of groups of carbon atoms were taken with a
TEM (Tunneling Electron Microscope) and an AFM (Atomic
Force Microscope).
Notice the scale bar (small black line in lower right of
picture). That line represents 2 nm (that’s 2 nanometers.)
If one nanometer equals 1x10-9 meters…
Then 62,500,000 of these pictures would fit in one meter!
That’s REALLY small!!
The Atom – Background & History
A summary of the models of the Atomic Jig-Saw Activity we did in class…
1. The Ancient Greeks Propose a Good Idea - Democritus:
a.
He was a Greek smart guy (a philosopher, a thinker) who suggested an indivisible form of
matter.
b.
He did no “science.” Testing of a good idea did not occur.
ATOMIC MODEL SKETCH:
2. Rediscovering a Good Idea: Dalton and Atoms (1766-1844):
a.
He proposed that all elements are composed of a “smallest particle” that could not be further
subdivided.
b.
He suggested that all atoms of the same element are identical.
c.
He noted that atoms of different elements can combine with each other in simple, whole number
ratios to form compounds. (Some simple compounds were known about by the early 1800’s.
Among the first was ammonia – NH3)
d.
Recall particle diagrams:
3
Chemical reactions (aka “chemical changes”) occur when atoms/molecules are:
1.) separated:
**Note that the atoms themselves
remain unchanged!**
2.) joined:
3.) rearranged:
ATOMIC MODEL SKETCH: (Think of a BILLIARD BALL!)
3. Discovery of Subatomic Particles…Dalton’s Undoing…
a.
b.
c.
Starting in the late 1800’s, particles smaller than atoms were being detected in experiments.
Eventually, this collection of particles was referred to as “sub-atomic” (smaller than atomic)
particles.
These particles are called electrons, protons and neutrons today.
Discovering Electrons:
Thomson (1897):
Experimented with Cathode Ray Tubes (CRTs)
e- 
+
A CRT is similar to your TV. It has an anode (A Negative electrODE) and a cathode (A positive electrode).
These are enclosed in an evacuated (air removed) glass container and when a charge is applied, the electrons flow
from anode to cathode through the open space of the glass container.
Thomson observed these particles and determined that the particles:
 Move at a very high speed (about 10% the speed of light)
 Have a negative charge
 Have a mass of about 1/2000 of a hydrogen atom (hydrogen is the smallest atom)
 Were the same regardless of which gas was used in the container or the metal used as the electrode
4
The “plum pudding model” attempted to explain atomic structure once the electron (e- or e) had been discovered.
An atom, according to this model, was a cluster of small positive and negative charges.
ATOMIC MODEL SKETCH: (Think of a CHOCOLATE CHIP MUFFIN!)
Discovering Protons – The Nucleus:
Rutherford (1911): Performed the Gold Foil Experiment
View the online simulations of Rutherford’s Gold Foil Experiment & make observations:
1. http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/ruther14.swf
2.http://www.physics.upenn.edu/courses/gladney/phys351/classes/Scattering/Rutherford_Scattering.html
a. Observation #1 –
Conclusion #1 –
**He called this the nucleus.
b. Observation #2 –
Conclusion #2 –
Real-life relation of “mostly empty space”:
If the nucleus is a marble/pea at midfield, the closest electron is at the goal post!
What Rutherford expected to see,
based on Thompson’s chocolate
chip muffin model

What Rutherford really saw.
Why did some alpha particles
bounce back? How did this change
the model of the atom? 
5
ATOMIC MODEL SKETCH: (Think of a PEACH!)
Discovering Neutrons:
Chadwick (1932):
a. Subatomic particle with no charge. (Hard to find without a charge!)
b. The nucleus is now considered to be composed of protons and neutrons.
c. This realization led to the “nuclear age” in the 1940s.
4. Bohr (Planetary) Model:
a. Rutherford proposed atoms had a positively charged nucleus, with negatively charged electrons orbiting
around it. The only problem: He could not explain why the electrons would not just spiral into the
nucleus (since opposite charges attract).
b. Bohr used data from atomic spectra to propose a soluton to the flaw in Rutherford’s model
c.Electrons must exist in fixed, stable orbits (or energy levels), which are at a specific distance from the nucleus.
ATOMIC MODEL SKETCH: (Think of an ONION!)
5. Modern Model: (Think of a FAN!)
a. aka - Quantum Mechanical Model, Wave Model, Wave Mechanical
Model, Cloud Model, Electron Cloud Model
b. There is a probability for finding the electron in regions of space
chemists today call “orbitals.”
6
Assignment #1 -Discoveries
1. What discoveries concerning the Modern Atomic Theory are each of the following scientists credited with?
a. Ernest Rutherford –
b. John Dalton –
c. J.J. Thomson –
d. Neils Bohr –
2. Draw each atomic model and identify which scientist is associated with it:
“Nuclear”
“Planetary”
“Billiard Ball”
“Plum Pudding”
Name of Scientist
Name of Scientist
Name of Scientist
Name of Scientist
Topic 1: Teachers Sign-Off: ___________________
7
**TOPIC 2**
Subatomic Particles & Symbols
You Should Know/Understand OR Be Able To Do:
1. Discuss the structure of an atom including location of the protons, electrons, and neutrons in relation to
the nucleus.
2. Describe the three basic subatomic particles (protons, electrons, and neutrons) in terms of relative
masses and electrical charges.
3. Use atomic number and mass number of an element to find the number of protons, electrons, and
neutrons in a particular atom.
Particle
Subatomic Particles & Their Properties
Relative
Approximate
Symbol
Location
Electrical
Relative Mass
Charge
(amu)

What does the unit “amu” mean?

Atoms are electrically
. This means that the number of
must equal the number of

What are “atomic numbers”?

What are “mass numbers”?
Actual Mass
(g)
.
Time for the Reference Tables….open up to the Periodic Table!
Examples of Writing Element Symbols (Notation):
8
Assignment #2: Atomic Number & Mass Number Worksheet
Complete the following chart and answer the questions below.
Element
Name
Atomic
Number
Number of
Protons
Number of
Neutrons
Number of
Electrons
Mass Number
carbon
12
8
8
hydrogen
1
6
hydrogen
14
2
nitrogen
14
1
92
2
146
cesium
82
11
12
47
tungsten
108
110
45
24
80
52
Silver
107
76
114
1. How are the atomic number and the number of protons related to each other?
2. How do the number of protons, number of neutrons, and the mass number relate to each other?
3. What is the one thing that determines the identity of an atom (that is, whether it is an oxygen atom or a carbon
atom, etc.)?
Topic 2: Teachers Sign-Off: ___________________
9
**TOPIC 3**
Subatomic Particles & Isotopes
You Should Know/Understand OR Be Able To Do:
1. Use atomic number and mass number of an element to find the number of protons, electrons, and
neutrons in a particular atom.
2. State how isotopes of an atom differ.
3. Interpret and write isotopic notation.
NOTES:
There are two different naturally occurring types of chlorine atoms, represented by Cl-35 and Cl-37.
 What is the difference between these two?!
Isotopes =
“Iso-” is a prefix that means…
So… What are three things that are the same between atoms that are isotopes?
What are two things that are different?
Remember:
1. The number of protons defines the element.
2. The number of neutrons determines which isotope of a given element you have.
ISOTOPES AND ATOMIC MASS
1. Atomic Mass: We need to differentiate between atomic mass and mass (number) of an atom!
Atomic Mass -
Mass Number -
10
2. Look at a Periodic Table....Atomic mass is given to a number of decimal places.
This is because, in most cases, there are a number of naturally occurring isotopes.
For example:
A natural sample of C (atomic mass = 12.011 amu) is a mixture of C-12 (98.89%) and
C-14 (1.11%).
A natural sample of N (atomic mass = 14.007) is a mixture of N-14 (99.63%) and
N-15 (0.37%).
Consider the following periodic table information for carbon:
[Notice on this symbol that the location of the atomic mass and the atomic number seem to be
reversed from how it shown on the Periodic Table in our reference tables. Be OK with that… the
atomic mass is always the number that is bigger in value, regardless of whether it is above or
below the symbol. European and Canadian charts show it this way.]
Carbon's atomic number is 6, has an average atomic mass of 12.011 amu, and carbon's most common isotope has
a mass number of 12 amu. Therefore, the most common type of carbon atom has 6 protons, 6 neutrons and 6
electrons. Another naturally-occurring isotope of carbon is C-14, but it is rare in comparison to the amount of C12 in nature.
Isotope
Hydrogen-1
Hydrogen-2
(deuterium)
Hydrogen-3
(tritium)
Atomic Number of Number of
Number Protons
Neutrons
Number of Mass Number
Electrons (amu)
1
1
0
1
1
1
1
1
1
2
1
1
2
1
3
Hydrogen has three isotopes, each with a whole number for a mass number.
Yet, hydrogen as an element has an atomic mass on the periodic table of 1.0079 amu.
Which isotope of hydrogen must be the most abundant, based on this data?
11
Assignment #3: Isotope Chart
Complete this chart to test your ability to work with isotope symbols.
Element
Name
Atom
(Isotope)
Atomic
Number
Number Number of
of Protons Neutrons
Number of
Electrons
Mass
Number
Nuclear
Charge
Overall
Atomic
Charge
Fe – 53
Aluminum27
19
18
86
222
90
38
Assignment #4: Atomic Structure Review
1. An atom of chlorine-35 contains how many:
a. protons?
b. neutrons?
c. electrons?
2. The three isotopes of oxygen, O-16, O-17 and O-18, all contain the same number of
and
.
How many are there?
3. The charge on the nucleus of a nitrogen atom is ____________. (Give both sign and number.)
4. Compared with a proton, an electron has _____________ mass.
(more or less)
5. Write the symbol for a possible formula of an isotope of
24
12
Mg . ___________
6. What particles contribute to the mass number of an atom? __________________________
12
7. The Rutherford experiment involved shooting alpha particles at ______________________.
8. Most of the alpha particles went through the foil undeflected. Rutherford concluded that:
9.A few of the alpha particles were deflected backward. Rutherford concluded that:
Topic 3: Teachers Sign-Off: ___________________
**TOPIC 4**
Calculating the Atomic Mass from Mass and Abundance of Isotopes
NOTES:
The atomic mass values listed on the periodic table are what are known as “weighted” averages of the naturally
occurring masses of the isotopes. You are somewhat familiar with the notion of a weighted average. Many of
your course grades are determined this way. If your teacher says to you “tests are worth 50% of your grade, HW
30%, etc” then this is an average calculated by giving each category a particular “weighting” or “percentage”
value in the calculation of the grade. For atomic mass calculations, the “weighting” given a particular isotope is
the “percent abundance” of that isotope in a natural sample.
Obviously, the isotope that is most prevalent in a natural sample of an element has the most effect on the
calculation of atomic mass for that element. For example, copper has 2 naturally occurring isotopes, Cu-63 and
Cu-65. The atomic mass of Cu is 63.546 amu. Therefore we may rightfully conclude that Cu-63 is the more
abundant isotope.
Other Resources:
View “Chem Guy” on You Tube: http://www.youtube.com/watch?v=BTA3OMuwhyI
Example:
Cr has 4 isotopes: Cr-50, 4.35% abundance, 49.946 amu; Cr-52, 83.79% abundance, 51.941 amu; Cr-53, 9.50%
abundance, 52.941 amu; Cr-54, 2.36% abundance, 53.939 amu. Calculate the average atomic mass.
49.946*amu(.0435) + 51.941amu*(.8379) + 52.941amu*(.0950) + 53.939amu*(.0236) = average amu
2.1727 amu + 43.521 amu + 5.0294 amu + 1.2730 amu = 51.996 amu
13
Assignment #5: Calculating Average Atomic Masses
Isotopes
Solve the following:
1. Boron has two naturally occurring isotopes: boron-10 (abundance = 19.8%) and boron-11 (abundance =
80.2%). Calculate the atomic mass of boron.
2. Calculate the atomic mass of magnesium. The three magnesium isotopes have atomic masses and relative
abundances of 23.985 amu (78.99%), 24.986 amu (10.00%), and 25.982 (11.01%).
3. Calculate the atomic mass of titanium. The five titanium isotopes have atomic masses and relative
abundances of 45.953 amu (8.0%), 46.952 amu (7.3%), 47.943 amu (73.8%), 48.948 (5.5%), and 49.945
amu (5.4%).
4. Copper has two naturally occurring isotopes: copper-63 (69.17%) and copper-65 (30.83%). Calculate the
average atomic mass of copper.
5. Chlorine has two naturally occurring isotopes: chlorine-35 (75.77%) and chlorine-37 (24.23%). Calculate
the average atomic mass of chlorine.
14
IONS
When neutral atoms lose or gain electrons they become ions.
If a neutral atom gains one or more electrons there is an overall increase in negative charges. Negative ions are
called anions.
Notice how the size of the fluoride ion is larger than the fluorine atom:
If a neutral atom loses one or more electrons there is an overall increase in positive charges. Positive ions are
called cations.
Notice how the calcium ion is smaller than the calcium atom:
Topic 4: Teachers Sign-Off: ___________________
15
**TOPIC 5**
Review Materials
Assignment # 6 Atomic Structure Drill
1. Find the number of p, n, and e- in the following species.
Ti2+
K+
N3-
Br
H+
H-
2. In the modern model of the atom, each atom is composed of three major subatomic (or fundamental) particles.
a) Name the subatomic particles contained in the nucleus of the atom.
b) State the charge associated with each type of subatomic particle contained in the nucleus
of the atom.
c) What is the net charge of the nucleus?
Answer the following multiple choice questions by circling the best response.
3. The number of neutrons in the nucleus of an atom can be determined by
(1) adding the atomic number to the mass number
(2) subtracting the atomic number from the mass number
(3) adding the mass number to the atomic mass
(4) subtracting the mass number from the atomic number
4. As an atom becomes an ion, its mass number
(1) decreases
(2) increases
(3) remains the same
5. The following equation represents the formation of a
(1) fluoride ion, which is smaller in radius than a fluorine atom
(2) fluorine atom, which is smaller in radius than a fluoride ion
(3) fluoride ion, which is larger in radius than a fluorine atom
(4) fluorine atom, which is larger is radius than a fluoride ion
6. Which statement best describes electrons?
(1) They are positive subatomic particles and are found in the nucleus.
(2) They are positive subatomic particles and are found surrounding the nucleus.
(3) They are negative subatomic particles and are found in the nucleus.
(4) They are negative subatomic particles and are found surrounding the nucleus.
7. In which list are the elements arranged in order of increasing atomic mass?
(1) Cl, K, Ar
(2) Fe, Co, Ni
(3) Te, I, Xe
(4) Ne, F, Na
8. The atomic number of an atom is always equal to the number of its
(1) protons, only (2) neutrons, only
(3) protons plus neutrons
(4) protons plus electrons
9. The nucleus of an atom of K-42 contains
(1) 19 protons and 23 neutrons (2) 19 protons and 42 neutrons
(3) 20 protons and 19 neutrons (4) 23 protons and 19 neutrons
16
10. What is the total number of electrons in a Cu+ ion?
(1) 28 (2) 29 (3) 30 (4) 36
11. After a neutral sulfur atom gains two electrons, what is the resulting charge of the ion?
12. Which particles are found in the nucleus of an atom?
(1) electrons, only
(2) neutrons, only
(3) protons and electrons
(4) protons and neutrons
13. What is the total number of neutrons in an atom of an element that has a mass number of 19 and an atomic
number of 9?
(1) 9
(2) 19
(3) 10
(4) 28
14. A neutral atom contains 12 neutrons and 11 electrons. The number of protons in this atom is
(1) 1
(2) 11
(3) 12
(4) 23
15. What is the total number of electrons in a Cr3+ ion?
(1) 18
(2) 21
(3) 24
(4) 27
16. Which statement is true about the charges assigned to an electron and a proton?
(1) Both an electron and a proton are positive.
(2) An electron is negative and a proton is positive.
(3) An electron is positive and a proton is negative.
(4) Both an electron and a proton are negative.
17. What is the charge of the nucleus in an atom of oxygen-17?
(1) 0
(2) -2
(3) +8
(4) +17
18. How many electrons are contained in an Au3+ ion?
(1) 76
(2) 79
(3) 82
(4) 197
Topic 5: Teachers Sign-Off: ___________________
Assignment #7: Unit Pre-Test
1. a. How do isotopes of the same element differ from each other?
b. If the two isotopes of Cl are Cl-35 and Cl-37, which predominates (which is there more of)?
Explain by using the atomic mass of Cl in your periodic table.
2. What is the charge on the nucleus of…(a) a Cr atom?
(b) a Ni atom?
(c) a sodium atom?
17
3. Fill in the following table:
Symbol
59
27
# protons
# neutrons
# electrons
atomic mass
Co
8
70
31
atomic #
10
16
Ga 3
17
20
18
4. Fill in the following table:
Particle
Charge
Mass
+1
1
Location
neutron
electron
5. Describe Rutherford’s experiment and his conclusions.
6. How many significant figures in the following numbers:
a. 2.080
b. 0.009
c. 1.009
d. 1800.
7. What is the correct answer (CORRECT NUMBER of SIG FIGS AND UNITS)?
a. The volume of a box which is 1.2 meters long, 0.23 meters wide and 10.6 meters high
18
b. The density of a piece of metal that weighs 3.456 grams and whose volume is 1.23 ml.
c. The area of a rectangle which is 0.0230 meters long and 1.0230 meters wide
8. Convert the following
a. 2.49 Liters to ______________mL
b. 18.6 mL to ____________Liters
c. 16 cm to ____________meters
9. Calculate the following:
a. Silver has 2 isotopes: Ag-107 has a mass of 106.905 amu and an abundance of 52.00%
and Ag-109 has a mass of 108.905 amu and an abundance of 48.00%. What is the atomic
mass of silver?
b. Gallium, which has an atomic mass of 69.723 amu, has 2 naturally occurring isotopes, Ga-69
and Ga-71. Which isotope occurs in greater abundance? Explain.
**REMEMBER TO MAKE SURE YOU KNOW/UNDERSTAND
OR ARE ABLE TO DO EACH OF THE OBJECTIVES GIVEN
AT THE TOP OF EACH TOPIC PAGE!**
19