Download TE 802 Group Unit Plan - Stephen Stauffer MATC Portfolio

TE 802 Group Unit Plan
------------------------------------------------------------------Names of Group Members AND their contributions:
1. Travis (c 4.8A, c 4.8B, c 4.8C)
2. Stephen (4.8 d, 4.8 g)
3. Kyle (4.8 e, 4.8 f)
4. Liz (4.8h, 4.8 i) (also did a stellar job helping out with the OPM table!)
5. Caitlin (4.9a, 4.9b, 4.9 c)
Topic: Periodic Table
Abstract
This unit addresses the periodic table and how it can be used to determine the subatomic particles
present in the atoms of each element, and what the properties of those particles are. Additionally
students will learn to use the periodic table to determine the oxidation states, electron
configuration, and electron orbitals for elements. Students should already have completed units on
the history of the atom and atomic theory. They should be able to identify the major experiments
that have led to the development of the current model of the atom from its modest beginnings as a
single indivisible piece of matter. The units following this will cover chemical bonding and formula
writing.
CLARIFYING YOUR GOALS
A. School Science Approach
Facts
-Protons and neutrons make up the nucleus of an atom, and it is surrounded by electrons.
-Periodic table groups elements based on chemical and physical properties, with certain atomic trends
appearing as you travel across a period or down a group:
-Atoms are mostly empty space; there is a small, dense nucleus at the center, then a large amount of space
that electrons “live” in.
-Electrons are located in clouds around the nucleus, and that “cloud” represents an area of very high
probability where they may be found. Nodes represent areas where electrons have a 0% chance of being
found; as you increase the principle quantum number, so too do you increase the number of nodes found
in any single orbital.
Vocabulary Words
Orbitals: Each of the actual or potential patterns of electron density that may be formed in an atom or
molecule by one or more electrons and that can be represented as a wave function
Protons: positively charged particles in the nucleus of the atom; characteristic particle of the element
Neutrons: neutrally charged particles in the nucleus of the atom; isotopes of the same element have a
different number of neutrons
Electrons: electrically charged particles surrounding the nucleus of the atom; ions of the same element
have a different number of electrons
Nucleus: The positive center of an atom that contains the majority of its mass
Periodic Table: A table of all known chemical elements arranged by mass, atomic properties, and electron
configuration.
Ionization Energy: Energy required to completely remove an electron from an atom
Physical change: Bonds remain unchanged, while the state or some other physical property (color, size,
and temperature) may change
Chemical change: Bonds are broken and formed between new atoms, changing the chemical composition
Metal:
● good conductors of heat and electricity
● high densities
● high melting points and boiling points
● generally form cations as they lose electrons
● located in the middle and far left of the periodic table
Metalloid: Generally share properties with metals and nonmetals without being classified one way or the
other. These include boron, silicon, germanium, arsenic, antimony, and tellurium, with some lists
occasionally including aluminum, carbon, selenium, polonium, and astanine.
Nonmetal:
●
●
●
poor conductors of heat and electricity
in solid form, they are dull and brittle
they have significantly lower melting points and boiling points
● non-metals have high electronegativity
● generally form anions as they gain electrons
● located on the far right of the periodic table [with the exception of hydrogen]
Electronegativity: The tendency of an atom to draw electrons in towards itself
Hund’s Rule: Every degenerate orbital must be filled with one electron before any orbital gets two; and
each electron in singly-occupied orbitals has the same spin.
Aufbau Principle: Electrons fill the lowest energy orbitals before moving to higher energy orbitals
Pauli Exclusion Principle: No electrons with the same spin may occupy an orbital at the same time.
Valence electrons: Electrons located in the highest energy orbital
Ions: Elements with differing charges; the different charge comes from gaining or losing electrons.
Isotopes: Elements with differing masses but identical protons/electrons; the extra mass comes from extra
neutrons in the nucleus.
Diagrams or standard representations
Thomson Model
Bohr Model
D Orbitals
Orbitals!
Lewis Dot Structure
Periodic Table
Formulas or problem-solving skills
Formal charge of an atom: Number of valence electrons - number of non-bonding electrons - [number of
bonded electrons / 2]
General understanding of atomic trends: It is important for students to understand the WHYs behind
atomic trends.
● Atomic radius increases going down a group because more orbitals are being added on; it
decreases going across a period because the attraction between the positive nucleus and the
electrons increase as the number of valence electrons increase
● Electronegativity increases going across a period because atoms need fewer electrons to reach a
full, stable octet; it decreases going down a group because more orbitals are being added on, and
they are shielding the positive charge of the nucleus
● Ionization energy increases going across a period because the atoms need fewer electrons to reach
a full, stable octet so they are less inclined to give up electrons; it decreases going down a group
because more orbitals are being added on, so they shield the positive charge of the nucleus,
allowing the electrons in the valence shell to pull away more easily.
B. Standards and Benchmarks
Benchmarks for Science Literacy
1. History often involves scientific and technological developments. 1C/H3b
2. From time to time, major shifts occur in the scientific view of how things work. More often, however,
the changes that take place in the body of scientific knowledge are small modifications of prior
knowledge. Continuity and change are persistent features of science. 1A/H2*
3. No matter how well one theory fits observations, a new theory might fit them just as well or better, or
might fit a wider range of observations. 1A/H3a
National Science Education Standards
http://www.nap.edu/openbook.php?record_id=4962&page=111
Michigan Content Expectations
C4.8A- Identify the location, relative mass, and charge for electrons, protons, and neutrons.
C4.8B- Describe the atom as mostly empty space with an extremely small, dense nucleus
consisting of the protons and neutrons and an electron cloud surrounding the nucleus.
C4.8C- Recognize that protons repel each other and that a strong force needs to be present to
keep the nucleus intact.
C4.8D- Give the number of electrons and protons present if the fluoride ion has a -1 charge.
C4.8e- Write the complete electron configuration of elements in the first four rows of the periodic
table.
C4.8f- Write kernel structures for main group elements.
C4.8g- Predict oxidation states and bonding capacity for main group elements using their
electron structure.
C4.8h- Describe the shape and orientation of s and p orbitals.
C4.8i- Describe the fact that the electron location cannot be exactly determined at any given
time.
C4.9A- Identify elements with similar chemical and physical properties using the periodic table.
C4.9b- Identify metals, non-metals, and metalloids using the periodic table.
C4.9c- Predict general trends in atomic radius, first ionization energy, and electonegativity of the
elements using the periodic table.
Synthesized Objectives
Students will be able to identify ions and use oxidation numbers in order to determine how
elements bond to one another. Students will also be able to appropriately determine how the
bonding in a compound changes as the oxidation numbers in the compound change. (C1.1f,
C4.8D, C4.8g)
Students will be able to group elements in the periodic table using chemical and physical
properties into metals, non-metals, and metalloids, and be able to predict the general trends of the
elements in the periodic table using data and empirical evidence. (C4.9A, C4.9b, C4.9c, C1.1D,
C1.1g)
Explain the structure of the atom, including the location and properties of its subatomic particles.
(4.8A, 4.8B, 4.8C)
Compare and contrast new and old atomic models using orbitals and probability. (4.8h, 4.8 i)
Write electron configurations for various elements and identify patterns of electron configuration
using the Periodic Table. (4.8 e, 4.8 f, C1.1d)
C. Big Ideas
Atoms are elementary particles that are the building blocks of everything around us. Their structure is
surprisingly simple. They are mostly empty space, with a tiny, extremely dense nucleus. Within this dense
nucleus are the important parts of the atom, the protons, and neutrons. Protons have a mass of 1.67262158
E-27 kg, and a positive charge therefore they repel each other very strongly so it is necessary for a very
strong force to be present to keep the nucleus intact. Neutrons have a mass of 1.674927 E-27 kg, and
carry no charge. Surrounding the nucleus there is a ‘cloud’ of electrons. Which is referred to as a cloud
because the exact location of an electron cannot ever be pinpointed exactly. This is because electrons
move around in order to help atoms of differing elements to bond with one another, also due to its
extremely small mass (9.10938188 E-31 kg). Electrons organize themselves into what are commonly
referred to as orbitals. Be careful with the word orbital they are not like planets that orbit; they are energy
levels that electrons occupy depending on their state of excitement. These orbitals come in several
different shapes, the most common shapes for lower energy elements are s and p orbitals. An s orbital is
sphere around the nucleus. A p orbital is shaped like an oblong 8 with the nucleus at the center of the 8
where the two ovals intersect. The orbitals are filled based off of several rules, first of which is the
Aufbau principle which says electrons want to occupy the lowest state of energy so they fill from the
lowest up. An electron will not enter a higher state of energy unless all the lower orbitals are filled. An
orbital is considered full when it contains two electrons of opposite and equal spin. (Pauli Exclusion
Principle) Electrons don’t like to be paired up if they can avoid it though and Hund’s Rule states that
electrons will try for the maximum amount of unpaired occupied orbitals. So all orbitals in a given energy
level fill with one electron before cycling back through adding the second electron then if more electrons
remain they start to fill the next energy level. What is an important thing to look for in this process is a
special type of electron called valence electrons, which means the outermost or highest energy electrons.
These electrons because of their high energy are the ones that will be the biggest factor in chemical
reactions. The Lewis dot structure is how you represent the number of valence electrons for an atom. The
Lewis dot structure is the elemental symbol surrounded by the number of valence electrons. The number
of valence electrons is determined by how many electrons reside in the highest energy level (not sublevel
so it would be all the electrons in for instance 6s and 6p not just 6p is 6 was the highest energy level
represented).
The periodic table is a table with all the known elements in it. They are not just thrown on there,
they are organized very specifically. They are organized so all the elements with similar chemical and
physical properties are near one another. There are many ‘groups’ on the periodic table such as metals,
non-metals, metalloids, and gases. The elements are ordered by increasing atomic number. There are
many trends in the periodic table because of the way they are grouped. Atomic radius of an element
increases as you go down the periodic table and from right to left. Ionization energy increases as you go
up the periodic table and as you go left to right across it. And electronegativity increases as you go up the
periodic table and from left to right.
D. Observations, Patterns, and Models
Observations or experiences
(examples, phenomena, data)
WHAT
Patterns
(laws, generalizations,
graphs, tables, categories)
HOW
Models
(explanations, theories)
-Protons have mass of 1.67262158
× 10-27 kg
-Neutrons have mass of 1.6749 x
10-27 kg
-Electrons have mass of
9.10938188 × 10-31 kg
-Carbon has a mass of 12.001 amu
-Hydrogen has a mass of 1.01 amu
-Oxygen has a mass of 15.999 amu
-Rutherford Gold Foil Experiment:
Dense nucleus, protons inside the
nucleus and electrons outside the
nucleus.
-Electron only found in certain
areas surrounding the nucleus (s
orbital, p orbital, d orbital, nodes)
-Different models of atom (Dalton,
Thomson, Bohr, Quantum)
-Atoms have mass
-Protons and neutrons have
more mass than electrons
-Nucleus is the most dense
part of the atom
-Generally, the higher the
atom is on the periodic table,
the more mass it has.
-The atomic number correlates
to the number of protons in an
atom.
-Atoms have a specific
structure
Atomic models:
Each atom has a characteristic
number of protons in a dense
nucleus with a corresponding
number of electrons to keep it
neutral. Atoms also have
neutrons that exist with the
protons inside the nucleus.
-Electrons fill orbitals in a certain
order.
-Each element has its own electron
configuration
-Electrons occupy space in the
electron cloud known as orbitals
which surround the nucleus.
-Aufbau Principle: Electrons
fill the lowest energy orbitals
before moving to higher
energy orbitals
-Pauli Exclusion Principle: No
electrons with the same spin
may occupy an orbital at the
same time.
-Hund’s Rule: Every
degenerate orbital must be
filled with one electron before
any orbital gets two; and each
electron in singly-occupied
orbitals has the same spin.
Electron configuration:
Every electron in an atom must
have a different quantum
number as no two may exist in
the same exact place at the same
exact time. Because of this,
electrons end up existing in
orbitals surrounding the
nucleus. The more electrons, the
more orbitals that are filled.
-Atoms that share
chemical/physical properties
are placed in the same groups
-Atoms are classified by
properties and periodic table
location
-Atomic radius decreases as
you go across a period,
increases down a group
-Electronegativity increases as
you go across a period,
decreases as you go down a
group
-Ionization energy increases as
you go across a period,
Periodic Table:
- Atomic radius: decreases
down a period due to a stronger
pull by the nucleus on the
electrons and increases down a
group due to an increase in
orbital shells
-Sodium reacts more violently than
lithium in water
-Potassium reacts more violently than
sodium in water
-Bromine and mercury are liquid at
room temperature
-Oxygen, nitrogen, hydrogen, helium,
etc are gases at room temperature
-Copper, carbon, magnesium, sodium,
lithium, etc are solids at room
temperature
-Alkali metals are in group 1
-Alkaline earth metals are in group 2
-Chalcogens are in group 16
-Halogens are in group 17
WHY
Electrons must be in the lowest
energy configuration for
maximum stability.
-Ionization energy:
increases down a period due to
atoms requiring fewer electrons
to reach an octet and decreases
down a group due to adding
more orbital shells which
increases shielding
-Noble gases are in group 18
-Atoms on the left side of the periodic
table are called metals
-Atoms on the right side of the table
are called nonmetals
-Atoms in between nonmetals and
metals are called metalloids
-Alkali metals/alkaline earth metals
react with halogens/chalcogens
-Fluorine is more electronegative than
oxygen
-Fluorine is more electronegative than
chlorine
-Fluorine has a higher ionization
energy than chlorine
-Fluorine has a higher ionization
energy than carbon
-Fluorine has a smaller radius than
oxygen
-Fluorine has a smaller radius than
chlorine
decreases as you go down a
group
-Electrons can be added or removed
-Pure elements have an oxidation
number (or formal charge) of 0, for
example, O(2), H(2), Mg(s). These
also have a neutral charge.
-Elements that lose electrons are
positive ions (cations):
● Mg2+ lost 2 electrons
● Na1+ lost 1 electron
● Al3+ lost 3 electrons
-Elements that gain electrons are
negative ions (anions):
● Cl1- gained 1 electron
● O2- gained 2 electrons
● F1- gained 1 electron
-Na radius is 180 pm
-Na1+ radius is 116 pm
-F radius is 50 pm
-F1- radius is 119 pm
-Mg radius is 150 pm
-Mg2+ radius is 86 pm
-O radius is 66pm
-O2- radius is 126 pm
-The more electrons an atom
gains, the more negative it
becomes
-The more electrons an atom
loses the more positive it
becomes.
-Cations are generally smaller
than neutral atoms
-Anions are generally larger
than neutral atoms
-Electronegativity:
increases down a period due to
atoms requiring fewer electrons
to reach a full octet and
decreases down a group due to
adding more orbital shells
which increases shielding and
decreases attraction
The table as we have it today
has been organized by
Mendeleev. The table
organizes atoms in rows
determined by the number of
protons and increasing number
of valence electrons, as well as
in columns determined by the
properties exhibited by the
elements. With the way these
elements are arranged, it allows
for observable trends to be
noticed along the table.
Ions and Valence Electrons:
Ions are when the number of
electrons in an atom does not equal
the number of protons. This
phenomenon gives atoms either a
positive or negative charge.
A certain amount of energy is
required to detach an electron from
an atom and so until that energy is
supplied (such as in a chemical
reaction) electrons generally
remain as a valence electron found
on an atom.
Atoms are most stable with 8
electrons in their valence shell, so
they either lose electrons if they
have less than 4 (metals/cations) or
gain electrons if they have more
than 4 (nonmetals/anions) to do
this.
Adding electrons increases the pull
exerted by the nucleus which
decreases atomic radius, and
removing decreases the pull which
increases atomic radius
Oxidation numbers involve the
transfer of electrons during
reactions. Free elements maintain
an oxidation number of zero, but
compounds contain oxidation
numbers that result in an overall
oxidation number of zero (unless
the compound itself has a charge
than the net number will match the
charge)
E. Students’ Prior Knowledge
A.
Accurate examples or ideas you can build on
- Know that protons, neutrons, and electrons are in atoms.
- Protons have a positive charge; electrons have a negative charge; neutrons are neutral.
- Protons and neutrons are in the nucleus; electrons orbit nucleus.
- Atoms can have a positive, negative, or neutral charge, depending on the number of protons,
electrons, and neutrons.
- Students are aware of elements, like Hydrogen, Carbon, Nitrogen, etc., but may not understand
what they are composed.
- Understand what atoms, molecules and compounds are.
B.
Common misconceptions
- The size of the proton, neutron and electron have the same mass.
- Neutrons contain a negative charge because neutron “sounds” negative.
- The atom is not mostly empty space, but rather it is solid.
- You can see the atom.
- Electrons are in the nucleus.
- Fill the electrons according to number and not by energy level.
- Confuse atomic number with valence number.
- The Periodic Table is just a table with no rhyme or reason for the placement of the elements.
ASSESSMENT AND ACTIVITIES
Synthesized Unit
Objectives
Summative
Assessment
Formative
Assessment
Major Activity
Compare and
contrast new and
old atomic models
using orbitals and
probability. (4.8h,
-Test questions.
1. Draw each of the
following: (a) 2) 1s
(b) 2px (c) 3dxy 2.
Draw a picture of
- Exit pass/bell work.
Have students draw
Bohr models for
different atoms, or
-Poster project. Students will
gather into groups of four and
draw atomic models
throughout the years
[Thomson --> Bohr -->
4.8 i)
Bohr’s model of the
oxygen atom and
make sure to label
(a) protons (b)
neutrons (c)
electrons. 3. What is
the majority of the
atom and which
experiment proved
this?
have them draw
different orbitals when
given the specific four
quantum numbers—
although that could
also be done just by
saying “draw the 3pz
orbital,” if that is better
for the students.
Identify ions and
use oxidation
numbers to
determine how
elements bond. Also
appropriately
determine how the
bonding in a
compound changes
as the oxidation
numbers in the
compound change.
Test / Quiz
1. List the number of
protons and
electrons found in
each listed ion:
a. H+
b. Li+
c. Zn2+
d. Al3+
e. NH(4)+
Bell Work What are valence
electrons and what
happens to an atom as
electrons are either
added to or taken away
from that atom.
2. Write about the
differences between
MgCl(4) and
MgCl(2). Make sure
to list the oxidation
numbers. It may be
helpful to draw the
Lewis dot structures
3. What are
oxidation numbers?
Current orbital model]. They
will need to say who
discovered the model, make
sure everything is labeled
properly, and explain in 2-4
sentences why the scientist
thought that was the proper
model, i.e. Rutherford’s gold
foil experiment. Hopefully
this will reach a variety of
kids, since it involves writing
and drawing, and it will help
them see how the model
evolved over the years, since
it’s not a direct lecture.
After students finish the bell
work, hold a class discussion
taking in ideas on what
students think happen to an
atom as electrons are either
added to or taken away from it.
Ask students what they think
an ion is giving the hint that it
Ticket to leave has to do with the number of
Label the oxidation
electrons on an atom. Guide
numbers in the reaction discussion to the idea that ions
CO(2) + H(2) -> CO + are atoms where electrons do
H2O. Mention how we not equal the number of
will be visiting these
protons on an atom.
types of reactions later
more in depth
Break up students into groups
and ask them to write up ways
to organize ions. Take
volunteers from groups to
share what they’ve come up
with. Afterwards, show them
how scientists describe and
label ions. For example: Cl-,
Ag+, Mg2+, Mg4+.
Hand out flashcards to the
groups that contain matching
scenarios such as the
following: “I have 18 electrons
but I have a net charge of
negative 1... What ion am I?” (
Cl-1) or “How many electrons
total does Mn+4 have?” Have
groups cycle through these
flash cards with a periodic
table to help them not only
become familiar with ions but
also with the table itself
Write electron
configurations for
various elements
and identify patterns
of electron
configuration using
the Periodic Table
Test / Quiz
1. Write the electron
configuration for the
following elements.
a. C
b. Cl
c. V
d. Se
e. Na
2. Write the electron
configuration using
noble gas
configuration.
a. Al
b. Ni
c. Br
d. Be
e. Nb
3. How are the
electron
configuration for
nitrogen and
phosphorous
similar?
4. What is Hund’s
Rule?
While during lecture
and homework
students will be asked
questions that have
them writing the
various electron
configurations for
different elements.
Furthermore, the
teacher will use cold
call to ask students to
recite the electron
configuration of
elements with a low
number of electrons.
Students will be asked
to explain the Aufbau
Principle, Pauli
Exclusion Principle
and Hund’s Rule.
Students will partake in an
activity that makes them fill an
apartment complex that
follows the rules of filling
electrons while writing
electron configurations.
Students will be told they are
an owner of a special
apartment building that has
special rules for filling its
occupants. The rules are as
follows:
1) The lowest floor must be
completely filled before the
next floor can be occupied.
(Aufbau Principle)
2) Nobody can have a
roommate until every room on
the floor has an occupant
(Pauli Exclusion Principle)
3) There can only be 2 people
per room.
4) There can be no male-male
or female-female roommates,
only male-female. (Hund’s
Rule)
After being told these rules
they will have to fill the
apartment by the rules.
After going over this, we will
discuss how filling the
apartment is like filling orbital
diagrams.
Then after the activity we will
go over how to do electron
configuration and how it
relates to orbital notation.
Once we have completed that
we will do several, and I mean
several, practice problems to
get them used to writing the
elements in electron
configuration.
Students will be
able to group
elements in the
periodic table using
chemical and
physical properties
into metals, nonmetals, and
metalloids, and be
able to predict the
general trends of
the elements in the
periodic table using
data and empirical
evidence. (C4.9A,
C4.9b, C4.9c,
C1.1D, C1.1g)
(questions modified
from Holt
Chemistry)
The difference
between the atomic
number of an element
in Period 2 and the
atomic number of the
element in Period 3
directly below it is?
The elements of a
group on the periodic
table have similar
what?
The alkali metals are
found in which group of
the periodic table?
The elements in Group
18 are called?
____________ is a
chemical property of
alkali metals
Which groups consists
of metals that typically
form ions with a 2+
charge?
The energy required to
remove the first
electron from a neutral
atom is called
____________ and
generally
____________ in
successive atoms
down through a group.
From top to bottom in
most groups of
elements, atomic radii
tend to?
Not including the noble
During lecture teach
will use cold call to
review the concepts
just covered
throughout lecture and
to review when lecture
moves on to multiple
days.
(borrowed/modified from
Holt Chemistry)
Have students observe
reactions of halide ions with
different reagents and analyze
data to determine
characteristics reactions of
each halide ion
Have students use references
to research the properties
different groups of elements.
Mix up the properties and have
students match them with the
different groups.
Make up some elements and
list some properties physical
and chemical and have the
students place them with the
group they would belong to.
gases, the elements
with the highest
electronegativities are
located in the
____________ of the
periodic table.
Explain the
structure of the
atom, including
the location and
properties of its
subatomic
particles. (4.8A,
4.8B, 4.8C)
Test
Sample questions:
1.) According to the
law of conservation
of mass, if Element
A has an atomic
mass of 2 mass units,
and element B has
an atomic mass of 3
mass units, what
is the mass of
compound AB?
2.) What number
uniquely identifies
an atom?
3.) What was
Rutherford able to
determine about the
structure of the atom
with
his gold foil
experiment?
4.) Which two
particles are found in
the nucleus of all
atoms but Hydrogen1?
5.) What is the
definition of an
atom?
Group presentation
on the important
experiments and
revisions of the
atomic
theory.
Constructing a Model
activity adapted from
Modern Chemistry by Davis,
Frey, Sarquis,
Sarquis.
Objective: The goal of this
For this activity,
activity is to have students
separate the class into construct a model of an
groups. Each group
unknown object by making
will be assigned
inferences about an object
a notable scientist
that is in a closed container
who helped advance
without touching or seeing
the atomic theory.
that object.
Examples include
Lavoisier
Procedure:
(conservation of
mass), Proust
1. Separate the class into
(definite
groups of 4-5 students. Each
proportions),
member of the group is
Rutherford (Gold
required to record the data
foil experiment), etc.
they collect and the
In class give each
inferences they made from
group time to
that data
research what the
2. Provide each group with a
scientist
sealed box filled with objects
did, and what
of different size,
inferences they drew mass, and shape.
from their
3. Without opening the box,
observations. Each
students must manipulate
group will then
the box in an attempt to
get 35-5 minutes to
determine which objects are
report their findings
in the box.
to the class. Have the 4. Record observations you
groups present on
made, such as weight,
number of objects, sound of
their scientist in
chronological order,
and after each
presentation discuss
how their
observations added to
the previous
understanding of the
atom.
objects, etc.
5. From those observations
try to determine what
objects are inside the box.
6. Open up the box and
compare your guesses to the
actual objects, which did
you get right, what did you
miss.
Discuss:
1. Real scientists use more
than one method to gather
data, how is this
illustrated by this activity?
2. Of the observations you
made, which were qualitative
and which were
quantitative?
3. What parallels can be
drawn between this activity,
and the experiments
conducted in order to
determine the structure of
the atom?