Download Christopher Warner Title: Element Project Educational Filters: The

Christopher Warner
Title: Element Project
Educational Filters:
The following factors determine the implementation of this unit.
o School District of Philadelphia Core Curriculum
o PA State Science Standards
o National Science Standards
o Former Student Feedback Regarding Applicability in High School Courses
Goals/Objectives:
- Explain the repeating pattern of chemical properties by using the repeating
patterns of atomic structure within the periodic table.
- Explain how structure determines function.
- Describe atoms as composed of even smaller sub-atomic structures.
- Name the parts of an atom.
- Describe the relationship between numbers of protons and neutrons and atomic
number.
- Explain how elements are arranged in the periodic table.
- Calculate atomic masses.
- Convert units.
- Use algebraic formulas to determine number of atoms in a sample.
PA Standards Covered:
• 3.1.10 Unifying Themes: Describe patterns of change in nature, physical, and
man-made systems.
•
3.2.10 Inquiry and Design: Apply knowledge and understanding about the nature
of scientific and technological knowledge.
•
3.4.10 Physical Science : Explain concepts about the structure and properties of
matter.
National Science Education Standards Covered
• Science Teaching Standard A
o Work together as colleagues within and across disciplines.
•
Science Inquiry Content Standard A
o Develop abilities to do scientific inquiry
o Understand scientific inquiry
•
Physical Science Content Standard B
o Develop an understanding of properties and changes of properties in
matter
Background Information
Early ideas of the atom began with Greek philosophy, but it was not until the
early 19th century that evidence of the existence of atoms began to be provided. Charles
Dalton proposed a hypothesis about the existence of atoms that was based on experiments
by the French chemist, Antoine Lavoisier. Lavoisier tested the idea that during chemical
reactions, matter cannot be created nor destroyed. This idea became the law of
conservation of mass. Another French chemist that Dalton considered was Joseph Proust,
who observed the idea that some substances (later called compounds) always have
elements in the same ratio by mass. This idea is now known as the law of definite
proportions (Brady, 1988).
Dalton explained these findings by hypothesizing that all matter is composed of
atoms, all atoms of the same element are identical, atoms of different atoms are different,
and that atoms unite in definite ratios to form compounds. These findings are all true
except the idea that all atoms of the same element are identical. Atoms have different
varieties called isotopes because different atoms of the same element can have a different
number of parts called neutrons (Brady, 1988).
Dalton also formed a law that suggests that the ratio of masses of one element that
combine with a constant mass of another element can be expressed in small whole
numbers. This law was not based on experimental data until another French chemist J.L.
Gay-Lussac worked with gas reactions. He observed that volumes of reacting gases were
in the ratio of small whole numbers. Amadeo Avogadro, an Italian physicist, supported
Gay-Lussac’s work with Dalton’s theory. Avogadro proposed equal volumes of gases,
under the same conditions, have the same number of molecules (Brady, 1988).
In the late 19th century the English scientist Michael Faraday found that chemical
changes occur when an electrical current is passed through certain chemical solutions.
These experiments were made using gas discharge tubes, glass tubes with a gas at low
pressure. J.J. Thomson discovered electrons experimenting with similar tubes called
cathode ray tubes, which are streams of electrons. He measured the bending of the path
of cathode rays and was then able to determine the ratio of the electron’s charge to its
mass. The proton was also discovered experimenting with cathode ray tubes. Rays
traveled in the direction opposite to that traveled by the cathode rays. These rays had an
opposite charge and had a mass 1800 times that of electrons (Smoot, 1987).
The English physicist Ernest Rutherford predicted the existence of a third type of
particle found in atoms. Walter Bothe and James Chadwick, who repeated Bothe’s work,
found high energy particles with no charge and a similar mass as the proton. This
particle is now known as neutrons (Smoot, 1987).
J.J. Thomson also noticed two kinds of neon atoms that were exactly alike
chemically, but different in mass. These different types of atoms of the same element are
called isotopes. Another English scientist, Henry Mosely, discovered the wavelength of
X-rays produced in an X-ray tube depends on the number of protons in the nucleus of the
atom. The number of protons was always the same for a specific element. This number
of protons, the atomic number, is the element identifier. The mass difference of isotopes
caused by the different numbers of neutrons in the nucleus, which determines the
particular isotope of the element (Smoot, 1987).
Ernest Rutherford and a younger Danish scientist, Niels Bohr, teamed up to
construct a model that incorporated their ideas. They made observations from an
experiment conducted by Hans Geiger and Ernest Marsden that involved sending
subatomic particles through a thin sheet of gold foil. Rutheford concluded that atoms
were composed of mostly empty space, a small, dense structure in the center called a
nucleus, and negatively charged particles called electrons traveling around the nucleus.
The Rutherford-Bohr model is commonly referred to as the planetary model because the
electrons orbit around the nucleus in the way that planets revolve around the sun (Smoot,
1987).
Through experiments involving substances that are exposed to energy, known as
spectroscopy, Bohr discovered that electrons travel around the nucleus in energy levels.
When electrons went to lower energy levels, light was emitted and could be studied.
Also, when electrons went to higher energy levels, light was absorbed in the same areas
in the spectrum and could be studied (Smoot, 1987).
Modern atomic theory states that electrons travel around the nucleus in electron
clouds where their location and speed cannot be determined. Only the probability of
where these electrons are can be calculated (Smoot, 1987).
Initial Activity for History Perspective
Students will be broken into groups. Each group will be assigned one of the
following atomic models: Dalton’s Model, Thomson’s Model, Rutherford’s Model,
Bohr’s Model, or the Modern Atomic Model.
They will each give power point presentations that answer the following questions:
1.
What did they find out about atoms?
2.
What experiments did they use to support their ideas?
3.
What ideas were correct or incorrect?
Activity A
Instructor should go through notes provided and incorporate exercises so that students
become comfortable with atomic number, mass number, atomic mass, and the parts of the
atom. Exercises and examples can be found at the end of this outline.
NOTES:
OBJECTIVE: Describe the size of an atom. Name the parts of the atom.
1) Size of atoms
a) How small are they?
i) Cannot be seen by normal microscope (smaller than microscopic)
ii) Diameter = 3 X 10-8 cm (of an aluminum atom)
iii) Size compared to object we know (tin foil example, p.88)
2) Parts of an atom
a) Mostly empty space
b) Nucleus
i) Found at the center of the atom
ii) Very small – size of a pinhead, if an atom was the size of a football stadium
iii) Very high density – lot of mass in a very small space (grape analogy p. 89, If a nucleus was
the size of a grape, it would weigh 18 billion pounds)
iv) Most of an atom’s mass is found in the nucleus;
v) Made of particles called protons and neutrons
(1) Protons
(a) Positively charged particles
(b) P for positive
(c) Protons have a charge of positive 1, charge of one proton = +1
(i) Charge of 5 protons = +5
(d) Have a mass of 1 amu (atomic mass unit)
(e) Why are amu’s used instead of grams? They’re easier to work with mathematically.
(i) Ex. Mass of 5 protons= 5 amu’s (atomic mass units)
(f) Sign of protons: p+
(2) Neutrons
(a) Particles with no charge, neutral
(b) N for neutral
(c) Have a mass of 1 amu
(d) Mass of 300 neutrons? 300 amu’s
c)
Outside the Nucleus
i) Electrons
(1) Negatively charged particles
(2) Each electron has a negative charge of negative1 (charge of one electron = -1)
(a) Charge of 5 electrons = -5
(3) Very low mass, negligible in calculating mass of atoms (electrons are not used to
calculate the mass of an atom)
(4) Sign/symbol of electron: e3) Different Atoms - 110 different elements, they all have different properties
a)
How is every element different?
i) Atomic Number
(1) Atoms of every different element have their own number of protons
(2) #of protons = atomic number
(3) # of protons is the sure way to identify an element.
(4) Exercise (look to p. 154)
(a) Atomic # of each, draw the box from the periodic table:
(b) Carbon - C
(c) Oxygen - O
(d) Gold – Au
(e) Mercury - Hg
(f) Your own example?
(i) Identify the element
1. #of protons =
a. 47 –
b. 28 –
c. 2 –
b) Isotopes
i) Isotopes are atoms of the same element.
ii) Isotopes have the same number of protons, but they do not have the same number of neutrons.
iii) Example: Hydrogen-1 has no neutrons. Hydrogen-2 has 1 neutron. They are atoms of the
same element (same atomic #), but they are not exactly the same.
iv) Mass Number
(1) The mass of a single atom.
(2) Mass number = # of protons + # of neutrons
mass = p + n
n = mass " p
(3) Mass numbers are used to name isotopes.
(a) Way they are named: “Name of element – mass #”
(b) Example: Fig.5 p. 92
(4) Exercise-mass numbers of different atoms; Figure out number of protons and neutrons in
each isotope:
(a) Carbon-14:
(b) Lithium-5:
(c) Helium-4:
!
c)
Atomic Mass
i) Weighted average of the masses of all the naturally occurring isotopes of that element
ii) What does this statement mean?
(1) Average mass of all the atoms of a single element found in the world.
iii) Math focus
(1) Calculate the atomic mass of boron, which occurs naturally as 20% boron-10 and 80%
boron-11.
OBJECTIVE: Explain why elements in a group often have similar properties.
1) What causes elements in a group to have similar properties?
a) The reason is the number of electrons in the atoms’ outer shell/level (valence electrons).
b) Why would these valence electrons cause similar properties?
i)
The number of electrons in the outer shell/level of an atom determines how it behaves and
how it can bond to other atoms.
2) Number of electrons allowed in each shell for the main group of elements:
a) 1st shell – 2 electrons
b) 2nd shell – 8 electrons
c) 3rd shell – 8 electron
3) Calculating the valence electrons of an element/atom
a) 1st – find out the number of electrons in the atom
i) If an atom is neutral (no charge), how many electrons are in the atom? The number of
electrons must equal the number of protons if the atom is neutral (no charge.)
b) 2nd – figure out how many electrons go into each shell
c)
3rd – count the amount of electrons in the outer shell
d) Examples of neutral atoms:
i) Helium
ii) Lithium
iii) Carbon
iv) Magnesium
v) Neon
vi) Oxygen
*Look at PT/answers
Assessment of Activities in Notes:
1. Number of protons given: What is the charge and the mass as a result of the
protons assuming the mass of a proton is 1 amu?
a. Example: 8 protons – charge is +8; mass is 8 amu
2. Number of neutrons given: What is the charge and the mass as a result of the
neutrons assuming the mass of a neutron is 1 amu?
a. Example: 5 neutrons – charge is 0; mass is 5 amu
3. Number of electrons given: What is the charge and the mass as a result of the
electrons assuming the mass of an electron to be 0 amu?
a. Example: 6 electrons – charge is -6; mass is 0 amu
4. Element given: Find the number of protons
a. Example: Helium – number of protons is two
5. Element given: Find the atomic number
a. Example: Lithium – atomic number is 3
6. Atomic number given: Find the element.
a. Example: Atomic number is 5 – Element is Boron
7. Number of protons given: Find the element.
a. Example: number of protons is 6 – Carbon
8. Isotope name given: Determine the mass of the atom.
a. Example: Carbon -12; mass = 12 amu
9. Isotope name given: Determine the number of protons, neutrons, and electrons if
the atom is neutral.
a. Example: Carbon-14; # of protons = 6, # of neutrons = 8,
# of electrons = 6
10. Mass percentages given: Determine the atomic mass of the element.
a. Example: If Boron-10 accounts for 20% of Boron found in nature, and
Boron-11 accounts for 80%, what is the atomic mass of Boron?
i. (.8)(11amu) + (.2)(10amu) = 10.8 amu
11. Element given: Determine the number of valence electrons assuming the atom is
neutral.
a. Example: Sodium - # of valence electrons = 1
Differentiated Instruction for Notes:
Make sure to assist students with learning difficulties and provide the typed out
steps above for solving each type of problem in the notes. Don’t provide answers unless
difficulties persist and give more problems to the student.
Activity B – Element Project
Each student will choose an element and complete a project regarding their element.
Artifacts:
Each student will draw, design, or make a model of an atom of their
element. Each student will also type up a report on their element that meet
a list of requirements.
-Atomic Model Aspect of Element Project
-Diagram of atom with correct number of protons, neutrons, electrons, and
valence electrons
-Calculations or shown work for number of protons, neutrons, electrons,
and valence electrons
-Research on element such as where it is found in nature and its uses
-List of some elements that bond to students’ sample
-Requirements for Typed Aspect Element Project
-Atomic Number
- Student must look at the periodic table to find the atomic number.
It’s the number found above the chemical symbol.
-Atomic Weight
- Student must research the percent weight of each isotope of their
element. Limit the number of isotopes to five, and make sure to
instruct the students that the percentages must sum up to 100%.
-Number of protons
-Number of protons = atomic number
-Number of neutrons
-Students must choose an isotope because different isotopes of the
same element have a different number of neutrons. # of neutrons =
mass number - # of protons.
-Number of electrons
-The number of electrons equals the number of protons if the
element is neutral.
-Number of valence electrons
-The number of valence electrons can be determined by using the
following instructions:
1st – find out the number of electrons in the atom
2nd – figure out how many electrons go into each shell
3rd– count the amount of electrons in the outer shell
-Calculation of number of atoms in a 5 gram sample
Example : Carbon
1x10"24 g
= 1.2x10"23 g
1amu
MassOfSample
5g
# ofAtoms =
=
= 5x10 23 CarbonAtoms
MassOfCarbonAtom 1.2x10"23
MassOfCarbonAtom(g) = 12amu *
!
-Uses of elements
This research would include where the atom is found and its
industrial uses.
-List of elements that can bond
-Compound Formulas
This is just an exercise to introduce students as to how to write
compounds.
Differentiated Instruction for Activity B
-Students with learning difficulties can simply do a research project on an element
and the common compounds that it is found in.
Activity C – Website Activity
Each student will be required to design a webpage that contains the information provided
in the Element Project (Activity B) including atomic number, atomic mass, melting point,
boiling point, where the element is found, the element’s uses in industry, and how the
element was discovered.
This part of the unit is cross-curricular in that students need to work in a computer lab.
Web design software is needed, and the assistance of a computer/technology instructor
may be necessary in assisting the students in designing their pages.
Activity D – Inquiry Activity
Have students conduct their own investigation using a scientific method where they test
for the physical and/or chemical properties of a Main Group element or a compound that
is seen in their everyday lives with that element present in the compound.
Inquiry requirement in conclusion: “Why do elements possess certain properties? Why
do other elements have similar properties?” The goal is to see whether students can
incorporate what they know about electrons and, in particular, valence electrons to
explain why groups of elements have similar properties. Time must be allotted for
students to discuss some of their research, their findings, and the periodic table. The
teacher should try to limit the guidance given to the students, however they should make
it available if the class as a whole struggles with the conclusion.
-
Some elements will not be practical if they cannot be found in safe sample, so
compounds can be used as long as it is explained why certain elements cannot be
tested and why compounds have different properties.
-
One of the main objectives is that students use a system or a method to figure out
some simple properties of substances such as density, state of matter, or melting
point. The method should include research, procedures, variables, observations,
and a conclusion that explains why elements and compounds have certain
properties.
Assessment of Activity D
REUBRIC
Title/Problem Statement
-Properties being diagnosed are introduced
and labeled.
Research
-Properties of elements vs. properties of
compounds are explained. Each property
being tested is defined and discussed in a
way that will lead to students making a
logical procedure to find out that property.
Procedure
-The procedure is logical and practical in
finding out the properties.
Variables
-Measurements and conditions that are not
the result of the procedure are labeled as
independent variables. The measurements
that are a result of the procedure are labeled
as dependent.
Observations
-The observations are listed and any work
required to calculate the property is shown.
Conclusion
-The conclusion explains WHY these
elements or compounds possess certain
properties. The student discusses how
structure determines function.
5 points
25 points
25 points
10 points
10 points
25 points
Bibliography
1. Brady, J. Holum. J. Fundamentals of Chemistry. 1988. John
Wiley and Sons. p.48-54.
2.
Smoot, R. Price, J. Smith, R. Chemistry: A Modern Course.
1987. p. 116-118, 120, 349.