Atomic History
Student Manual
Ms. Flint’s Chemistry Class
Dickinson High School
2013 – 2014
Getting Started: Is Air Made up of Matter?
With a partner answer this question in your notebook and give a brief explanation.
Lab Activity 1: What is Matter
1. Complete the What is Matter Lab provided at your lab station.
2. As you proceed with your lab activity, remember to Demonstrate safe lab practices.
 Read and follow all instructions on the station cards carefully.
 Make careful observations in your notebook.
 Ask questions.
 Discuss possible explanations.
 Leave your area clean and organized.
Reading to Learn 1: Matter
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Read the following passage carefully.
Take notes over key points in your notebook as you read.
Chemistry is the study of matter and the changes it undergoes. Matter is anything that has mass and has
volume. Mass is a measure of how much matter something has. Volume is a measure of how much space
something takes up.
You don’t have to be able to see something for it to qualify as matter. Matter can be in the form of a solid,
liquid or gas and can be a pure substance or a mixture. We will start our study of matter by looking at the
simplest form of matter—the atom.
In your notebook Record the definition of all key terms.
Reading to Learn 2: Early Atomic Theory
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Read the following passage carefully.
Take notes over key points in your notebook as you read.
The Greeks and the Atom
The idea that matter consists of small particles is as old as the ancient Greeks. Democritus (460 -370 BC), also
known as the Laughing Philosopher, suggested that all materials were made of tiny particles called atoms. The
word atom comes from the Greek word atomos meaning “indivisible” or “can not be divided” .
He believed that material things consist of an infinite number of very small atoms. He believed that atoms
were infinite in number, various in size and shape, and completely solid. Atoms were thought to move around
in a void repelling one another when they collided, or combining into clusters by means of tiny hooks and
barbs on their surfaces.
Other than changing places to form new clusters, atoms were thought to be unchangeable, ungenerated, and
indestructible. All visible objects were thought to be created from clusters of atoms. Any change in an object
was thought to be caused by atoms changing position. All interactions between atoms were believed to be
mechanical in nature, and to occur much like Velcro sticking together.
But, the Greeks never performed any experiments to test their ideas and a “non-atomic” view of material
won out over the Democritus’ idea. So the development of the atomic model went undiscovered for more
than 2000 years until the modern scientific revolution.
“By convention there is sweet, by convention there is bitterness,
by convention hot and cold, by convention color; but in reality
there are only atoms and the void.” - Democritus
The Solid Sphere Model
By the early 1800’s, the Law of Conservation of Mass, the Law of Definite Composition and the Law of Multiple
Proportions were commonly accepted ideas concerning the behavior of matter.
 Law of Conservation of Mass states that matter cannot be created or destroyed.
 Law of Definite Composition states that in any sample of a compound the masses of the elements are
in the same ratio.
 Law of Multiple Proportions states that when two elements combine to form more than one
compound, the fixed amounts of one element will combine with the other in a ratio of small whole
numbers
Influenced by these concepts, John Dalton took these ideas and the work of others,
including Democritus, and developed the first atomic model based on evidence collected
through scientific experimentation. He outlined his theory in what is known today as
Dalton’s Postulates.
Dalton’s Postulates
1. All matter consists of tiny, indivisible particles called atoms.
2. A chemical reaction is the result of rearrangement, combination, or separation of atoms. Atoms
cannot be created, destroyed or changed into different types of atoms in a chemical reaction.
3. Atoms of the same element are identical and have the same size, mass and chemical properties.
Elements have different properties as a result of differences in the atoms of each type of elements.
4. The atoms in a compound will combine in a definite, simple, whole number ratio each time.
The key difference between Dalton’s atomic theory and that of early Greek philosophers, such as Democritus,
was that Dalton based his model on scientific evidence and not just thought processes and beliefs.
“We might as well attempt to introduce a new planet into the
solar system, or to annihilate one already in existence, as to
create or destroy a particle of hydrogen.” – John Dalton
In your notebook Define all key terms.
 Record Dalton’s Postulates.
RTL Activity 1: Understanding Dalton’s Postulates
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Work with your partner to organize the cards provided.
Once complete, ask for the teacher to check that you have the correct arrangement of cards.
In your notebook, create a chart to organize the information from the cards.
Record all of the information from the cards accurately.
Put cards away neatly.
Reading to Learn 3: Thomson’s Plum Pudding Model
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Read the following passage carefully.
Take notes over key points in your notebook as you read.
Plum-Pudding Model
In 1859 a German physicist developed a vacuum tube that allowed a current to flow across a gap from the
cathode (-) to the anode (+). This was a type of radiation called cathode rays.
Later a different physicist suggested that cathode rays were actually streams of particles flowing from the
cathode to the anode, and these particles had the same properties regardless of the gas used to produce the
cathode rays.
In 1897, English physicist, J. J. Thomson performed a set of experiments that showed cathode rays could be
deflected away from a negative charge by both electric and magnetic fields. This discovery led Thomson to
suggest that cathode rays are negatively charged particles which he called electrons. He also determined that
the charge of an electron was about -1 x 10-19 coulombs.
Later, American physicist Robert Millikan performed his famous oil drop experiment to determine the exact
charge and mass of the electron.
The electron was no longer an idea but a real particle, but its mass is 1836 times smaller than the hydrogen
ion, the smallest known ion. So this suggested that the electron is just a small part of the atom. The term
subatomic particle was developed to describe these new particles that made up the atom.
Dalton's theory of an indivisible particle was no longer acceptable.
A theory that allowed for subatomic particles was needed. Thomson proposed that the
atom was a mass of positive charge with negative electrons embedded into it like raisins
in plum pudding, so it was dubbed the Plum Pudding model. (Think of a plum pudding
as being similar to a ball of chocolate chip cookie dough—the electrons are the
chocolate chips and the positive charge is the dough ball).
The table below outlines Thomson’s Cathode Ray experiments.
Soon after, Philipp Lenard noted that cathode rays could pass through very thin pieces of matter. So he
suggested that the atom was not a solid mass but mostly space. However, he did not have sufficient scientific
evidence to prove this.
I was told long afterwards by a distinguished physicist who
had been present at my lecture that he thought I had been
pulling their leg -J. J. Thomson
In your notebook Define all key terms.
 Create your own table to summarize the outline of Thompson’s experiments.
 Sketch a diagram of the Plum Pudding model of the atom.
 Which of Dalton’s Postulates was proven to be incorrect by Thomson’s experiments?
Lab Activity 2: Gold Foil Demo
As you observe the demonstration,  Carefully record your observations in your notebook.
 Record any questions you may have in your notebook.
 Think about possible explanations.
Reading to Learn 4: Rutherford’s Gold Foil Experiment
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Read the following passage carefully.
Take notes over key points in your notebook as you read.
Nuclear Atom Model
Ernest Rutherford was J.J. Thomson's student and decided to prove that the Plum-Pudding model was correct
and Phillip Lenard was wrong. So, Rutherford put the Plum-Pudding model to the test. In 1909, he performed
the famous Gold Foil experiment that led to the discoveries of the atomic nucleus and that the atom was
mostly space.
Rutherford set up an experiment to determine the effect of bombarding a thin sheet of gold foil with rays
from a radioactive source such as radium. He cut a small hole in a lead box so that the stream of rays could be
focused. Then he placed a thin sheet of gold foil and a viewing screen on the other side of the hole as shown in
the diagram.
The flashes of light produced when alpha particles strike the zinc sulfide screen can be observed through a
movable microscope. Most of the alpha particles passed through the gold foil undeflected, but a few were
deflected, and some even bounced back toward the source.
This was the expected
results for the Gold Foil
Experiment
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These were the actual
results for the Gold
Foil Experiment
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The Gold Foil experiment demonstrated that the mass of the atom was the same as predicted by Thompson's
model, but the volume of the mass was much smaller and seemed to be located in the center of the atom.
This dense center, known today as the atomic nucleus, has a volume one trillionth the volume of the whole
atom. The atomic atomic nucleus is a very small, very dense, centrally located, positively charged region in
the atom.
Thompson's Plum Pudding model was no longer workable. So, in 1911, Rutherford proposed the Nuclear
Atomic Model. Rutherford's new model stated that the atom was mostly empty space with a small, very dense
positively charged nucleus surrounded by negatively charged electrons located in the atomic space.
Soon after, the second subatomic particle was described as the proton. The proton is a positively charged
particle that is much larger than the electron and accounts for much of the mass in an atom.
It was almost as incredible as if you fired a 15-inch shell at a piece
of tissue paper and it came back and hit you. –Ernest Rutherford
In your notebook Define all key terms.
 Sketch a diagram of the Gold Foil Experiment.
 Answer the following questions:
1. Were the results of Rutherford’s experiment what he had expected? Write a brief explanation to
support your answer.
2. What was the major new development in atomic theory caused by Rutherford’s experiments?
Lab Activity 3: Atomic Target Practice Lab
1. Complete the Atomic Target Practice Lab provided at your lab station.
2. As you proceed with your lab activity, remember to Demonstrate safe lab practices.
 Read and follow all instructions on the station cards carefully.
 Make careful observations in your notebook.
 Ask questions.
 Discuss possible explanations.
 Leave your area clean and organized.
Reading to Learn 5: Modern Atomic Theory
Part 1
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Read the following passage carefully.
Take notes over key points in your notebook as you read.
The Planetary Model
According to Newtonian physics, if a negatively charged particle (electron) moves about a positively charged
particle (nucleus) energy will be released. This means that according to Rutherford’s model, the electron
should lose energy, slow down and move closer to the nucleus until it collides with the nucleus and destroys
the atom.
But, this doesn't happen and the Nuclear model can’t account for this phenomenon. So, a new model was
needed.
In 1913 Danish physicist, Neils Bohr, suggested applying the new ideas of the quantum theory to the Nuclear
model to explain why electrons don't spiral into the nucleus, and that a mathematical equation could show
the location of the electrons with respect to the nucleus.
Bohr's new model proposed that electrons are in fixed energy levels (n) he called orbits. These orbits are
arranged in concentric circles around the nucleus much like the planets orbit around the sun. The energy of
these orbits is quantized (a fixed amount) and electrons must absorb or release energy (photons) at
certain frequencies to move between energy levels.
This diagram represents Bohr’s
planetary model of the atom. Each
circle represents an energy level or
orbit.
The lowest energy level of an electron is the ground state and the higher energy levels are excited states. If an
electron absorbs a photon of a certain wavelength, the electron would be excited to a higher energy level
proportional to that wavelength. But, the lowest energy level possible would be the most stable so the
electron would immediately release the photons enabling the electron to return to its ground state.
If the quantum theory is true, particles that radiate energy must radiate energy as whole photons (quanta)
only. There could be no in between energy levels for the electron. It would have to jump from energy level to
energy level completely or not at all.
In the diagram below, the turtle on the stairs is behaving like an electron and the stairs represent the energy levels.
When the electron returned to the ground state this would produce a line spectrum instead of a continuous
spectrum.
The top picture in the diagram below shows a continuous spectrum while the bottom picture shows a line spectrum.
Lab Activity 4: Light Emission Demo
As you observe the demonstration,  Carefully record your observations in your notebook.
 Think about possible explanations.
 Did all of the mints emit light? Write a brief explanation for your observations.
Part 2
The Quantum Mechanical Model
In 1926, Austrian physicist Erwin Schrödinger used theoretical calculations and
experimental results to create and solve a mathematical equation that described the
behavior of the electron in a hydrogen atom. The modern description of electrons in an
atom comes from this equation.
In 1927, Werner Heisenberg proposed from a purely theoretical view that it was impossible
to know both the position and momentum of the electron simultaneously. This idea, which
became known as the Heisenberg uncertainty principle, dealt with the probability of an electron's position in
a region of space rather than its exact position.
In view of the uncertainty principle, it was suggested that the solutions to the Schrödinger wave equation be
taken as a description of the probability of finding electrons in certain areas of space. These solutions in the
form of numbers are called quantum numbers.
Quantum numbers not only describe specific quantized energy states for the electron but also a set of
probabilities for the position of the electron in a given energy level.
These probable positions known as atomic orbitals refer to a region in space where there is a high probability
of finding an electron, whereas an orbit is a definite path in space. Since the orbital does not have definite
boundaries, it is sometimes referred to as an electron cloud. The four quantum numbers can be used to
describe the probable location of each electron. More importantly, quantum numbers can describe the
electron configurations (arrangements) in all atoms.
The Schrödinger equation, the Heisenberg Uncertainty Principle and the quantum numbers produced a purely
mathematical view of the atom and the arrangement of the electrons. Like the Bohr model, the quantum
mechanical model restricts the energy of electrons to certain energy levels.
Unlike the Bohr model, the quantum mechanical model does not involve an exact path that the electron takes
around the nucleus. Instead, the location of each electron is described by the probability of finding the
electron in various locations. The probability of finding an electron within a certain volume of space around
the nucleus can be represented as a fuzzy cloud. Each of these regions are called orbitals.
In your notebook Define all key terms.
 Sketch a diagram of Bohr’s Planetary Model and of Schrödinger’s Quantum Mechanical Model .
As it so often happens, James Chadwick is once again represented by a footnote in Atomic History.
His first (and really only) claim to fame is that in 1932 he discovered the neutron. With its discovery
we have now discussed all three of the common subatomic particles which make up the modern
atom.
Oh, the cartoon character Jimmy Neutron is also named after him!
(Jimmy is often a nickname for James)
Atomic Theory Timeline
Major Developments in the History of the Atom
Applying What you Learned: Create an Atomic History Card Sort
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Using the template provided create a set of 25 cards containing key facts about the history and
development of the modern atomic theory.
For each of the five main scientists who made major contributions to atomic theory make five cards
that contain the following information:
1. A diagram of their atomic model (with the name of the model)
2. The scientist’s full name (a sketch would be a nice touch)
3. The date of their discovery
4. The major contribution or development made (discovered electron, described energy levels, etc. )
5. A summary of important information.
To receive full credit:
 Your writing needs to be neat and legible.
 If you make a mistake do not scribble. Just draw one line through the mistake.
 Your diagrams should be detailed, neat and include color and labels.
 Your cards are cut apart neatly and stored properly.
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When you have finished creating the cards, carefully cut them out and place them in the baggie
provided.
Write your name, the date, class period and the title “ Atomic History Card Sort” on the baggie in
sharpie marker.
Turn your sealed baggie into the Turn in Box.
!!! This baggie goes in your binder as part of your notes when returned to you. !!!
Check Your Understanding
Use your knowledge of atomic theory to answer the individual assessment. Pick this up from the teacher after
turning your folder in.