Name___________________________
Per. _____
Class Notes: Atomic Structure, Electrons, and the Periodic Table (Ch. 4-5)
NOTE: This set of class notes is not complete. We will be filling in information during class. If you are absent, it is your
responsibility to get the missing information from a fellow classmate.
I. Atom: smallest particle of an element that retains the properties of that element
II. Atomic Theorists
Democritus (460-370 BC)
Matter is made of tiny, solid, indivisible particles which he called atoms (from atomos, the Greek word
for indivisible).
Different kinds of atoms have different sizes and shapes.
Different properties of matter are due to the differences in size, shape, and movement of atoms.
Democritus’ ideas, though correct, were widely rejected by his peers, most notably Aristotle (384322 BC). Aristotle was a very influential Greek philosopher who had a different view of matter. He
believed that everything was composed of the four elements earth, air, fire, and water. Because at
that time in history, Democritus’ ideas about the atom could not be tested experimentally, the opinions
of well-known Aristotle won out. Democritus’ ideas were not revived until John Dalton developed his
atomic theory in the 19th century!
John Dalton (1766-1844)
1. All matter is composed of extremely small particles called atoms.
2. All atoms of one element are identical.
3. Atoms of a given element are different from those of any other element.
4. Atoms of one element combine with atoms of another element to form compounds.
5. Atoms are indivisible. In addition, they cannot be created or destroyed, just rearranged.
Dalton’s theory was of critical importance. He was able to support his ideas through experimentation,
and his work revolutionized scientists’ concept of matter and its smallest building block, the atom.
Dalton’s theory has two flaws:
a) In point #2, this is not completely true. Isotopes of a given element are not totally identical;
they differ in the number of neutrons. Scientists did not at this time know about isotopes.
b) In point #5, atoms are not indivisible. Atoms are made of even smaller particles (protons,
neutrons, electrons). Atoms can be broken down, but only in a nuclear reaction, which Dalton
was unfamiliar with.
III. Discovery of subatomic particles and refinement of the atomic model
JJ Thomson (1856-1940) discovered the electron, and determined that it had a negative charge, by
experimentation with cathode ray tubes. A cathode ray tube is a glass tube in which electrons flow
due to opposing charges at each end. Televisions and computer monitors contain cathode ray tubes.
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Thomson developed a model of the atom called the plum pudding model. It showed evenly
distributed negative electrons in a uniform positive cage.
Diagram of plum pudding model:
Ernest Rutherford (1871-1937)
Discovered the nucleus of the atom in his famous Gold Foil Experiment. In the experiment, alpha
particles (helium nuclei) produced from the radioactive decay of polonium streamed toward a sheet of
gold foil. To Rutherford’s great surprise, some of the alpha particles bounced off of the gold foil. This
meant that they were hitting a dense, relatively large object, which Rutherford called the nucleus.
Diagram of experimental setup:
Rutherford's Experiment
Questions about Rutherford’s experiment:
I. If gold atoms were solid spheres stacked together with no space between them, what would you
expect would happen to particles shot at them. Explain your reasoning.
2. What does Ernest Rutherford’s experiment suggest about the structure of the atom; in other
words, how can Rutherford’s evidence be used to correct the plum pudding model? Draw a
diagram.
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3. Can you explain why Rutherford concluded that the mass of the gold nucleus must be much
greater than the mass of an alpha particle (helium nucleus)?
4. Do you think that, in Rutherford's experiment, the electrons in the gold atoms would deflect the
alpha particles significantly? Why or why not?
5. Rutherford experimented with many kinds of metal foil as the target. The results were always
similar. Why was it important to do this?
Rutherford then discovered the proton, and next, working with a colleague, James Chadwick (18911974), he discovered the neutron as well.
Niehls Bohr
Developed the Bohr model of the atom (1913) in which electrons are restricted to specific energies
and follow paths called orbits a fixed distance from the nucleus. This is similar to the way the planets
orbit the sun. However, electrons do not have neat orbits like the planets.
Diagram of Bohr model:
Quantum Mechanical Model
This is the current model of the atom. We now know that electrons exist in regions of space around
the nucleus, but their paths cannot be predicted. The electron’s motion is random and we can only
talk about the probability of an electron being in a certain region.
Table: Subatomic Particles
Particle
Symbol
Charge
Mass (grams)
Electron
e-1
9.11 x 10 – 28 g*
Proton
p+
+1
1.673 x 10 – 24 g
0
Neutron
n
0
1.675 x 10 – 24 g
* The electron has only 1/1840 the mass of a proton or neutron.
Where found
In electron cloud
Nucleus
Nucleus
If this positive particle were the proton in a hydrogen atom, it would take a screen 1 mile across to
display the electron's orbit.
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A neutron walked into a bar and asked how much for a drink. The bartender replied, “For you, no
charge.”
If an electron weighed the same as a dime, a proton would weigh the same as a gallon of milk.
Although protons and electrons have such a different mass, they have the same magnitude of
charge:
The charge on a proton is
+1
The charge on an electron is
-1
Atomic Number
The periodic table is organized in order of increasing atomic number. The atomic number is the
whole number that is unique for each element on the periodic table. The atomic number defines the
element. For example, if the atomic number is 6, the element is carbon. If the atomic number is not
6, the element is not carbon.
The atomic number represents:
 the number of protons in one atom of that element
 the number of electrons in one atom of that element (with an ion, the electrons will be different)
**Therefore, protons = electrons in a neutral atom**
Isotopes
Atoms of an element with the same number of protons but different numbers of neutrons.
Two types of notation for isotopes:
Examples: finding protons, neutrons, and electrons for an element
Atomic Mass
Represents the average atomic mass of all isotopes of an element. Atomic mass is measured in
a.m.u. (atomic mass units). 1 amu = 1.660538 x 10 – 24 grams.
The masses of all elements are compared to carbon-12, which is defined as having a mass of exactly
12 amu.
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Atomic mass is the number on the periodic table that has several decimal places.
Determining Average Atomic Mass – use a weighted average technique
Example: Finding average atomic mass
Neon has 3 isotopes: Neon-20 has a mass of 19.992 amu and an abundance of 90.51%. Neon-21 has a
mass of 20.994 amu and an abundance of 0.27%. Neon-22 has a mass of 21.991 amu and an abundance of
9.22%. What is the average atomic mass of neon?
Electrons in Atoms
Electrons are found outside the nucleus, in a region of space called the electron cloud.
Electrons are organized in energy levels of positive integer value (n = 1, 2, 3,...).
Within each energy level are energy sublevels, designated by a letter: s, p, d, or f.
Each sublevel corresponds to a certain electron cloud shape, called an atomic orbital.
Analogy:
The electron cloud is like an apartment building.
The energy levels are like floors in the apartment building.
The sublevels are like apartments on a floor of the building. Just like there are different sizes of
sublevels, there are different sizes of apartments: 1 bedroom, 2 bedroom, etc.
The orbitals are like rooms within an apartment.
The electrons are like people living in the rooms.
What do these atomic orbitals look like?
Some examples:
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The shapes of the f orbitals are very complex and exotic. If you wish to view them, see this website:
http://antoine.frostburg.edu/chem/senese/101/electrons/faq/f-orbital-shapes.shtml
Each orbital can hold a maximum of 2 electrons.
An “s” sublevel contains 1 s orbital. How many total electrons can fit in an s sublevel? ___
A “p” sublevel contains 3 p orbitals. How many total electrons can fit in a p sublevel? ___
A “d” sublevel contains 5 d orbitals. How many total electrons can fit in a d sublevel? ___
An “f” sublevel contains 7 f orbitals. How many total electrons can fit in an f sublevel? ___
Electron Configurations and Orbital Diagrams
Three rules govern the filling of atomic orbitals.
1. The Aufbau Principle: Electrons enter orbitals of lowest energy first. The Aufbau order lists the
orbitals from lowest to highest energy: (“Aufbau” is from the German verb aufbauen: to build up)
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10 6p6 7s2 5f14 6d10
2. The Pauli Exclusion Principle: An atomic orbital may hold at most 2 electrons, and they must have
opposite spins (called paired spins).
3. Hund’s Rule: When electrons occupy orbitals of equal energy (such as three p orbitals), one
electron enters each orbital until all the orbitals contain one electron with spins parallel. Second
electrons then add to each orbital so that their spins are paired with the first electron in the orbital.
An electron configuration uses the Aufbau order to show how electrons are distributed within the
atomic orbitals.
How to read a segment of an electron configuration: Example
3p6
Examples of electron configurations:
Element
Total # of Electrons
Electron Configuration
Carbon
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Fluorine
Magnesium
Argon
Orbital diagrams show with arrow notation how the electrons are arranged in atomic orbitals for a
given element.
Element
Orbital Notation
Carbon
Fluorine
Magnesium
Argon
Valence electrons: Electrons in the outer energy level of an atom. They are like the front lines of an
army, because they are the ones involved in chemical reactions (valence electrons get shared or
transferred during reactions).
The number of valence electrons that an atom has is directly responsible for the atom’s chemical
behavior and reactivity.
How to determine valence electrons from electron configurations: Remember they are the electrons
in the outer energy level.
Element
Electron Configuration
# Valence
Electrons
Li
Be
B
C
N
O
F
Ne
More about the structure of the periodic table:
Group (or Family): a vertical column
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Electron Dot
Structure
Period: a horizontal row
Labeling on the periodic table: (you should have your s, p, d, and f blocks colored as well)
 Number of valence electrons for all representative element groups (Group A elements or
Groups 1-2 and 13-18)
 Alkali Metals: most reactive metals
 Alkaline Earth Metals
 Transition Metals: Group B elements or Groups 3-12
 Halogens: most reactive nonmetals
 Noble Gases: unreactive because they have 8 valence electrons
 Lanthanide Series (inner transition metals)
 Actinide Series (inner transition metals)
Notice that the elements in a family all have the same number of valence electrons (example: Alkali
Metals all have 1 valence electron). Therefore, these elements will behave in similar ways.
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