-The invention of the scanning electron
microscope has allowed evidence about the
atom to be gathered directly
-Atomic models give information about
chemical bonding, chemical reactions,
chemical and physical properties
-Greek philosopher Democritus was first to state that
matter was made of separate, distinct particles. These
particles could not be divided without losing its
properties; called these particles “atomos”
-Atom now represents the smallest particle capable of
chemical interactions
-In the 17th century, science (experiments) was used to
prove that matter is composed of atoms
-Observed that compounds always contain a set ratio
of elements and these ratios are unique: law of definite
composition
-In 1803, John Dalton created an atomic model based
on experimental evidence rather than philosophy
-Based on two premises: 1. combinations of atoms from
different elements form different compounds; 2. atoms
of different elements have different masses
-Dalton assigned masses to atoms so he determine their
relation to each other
-The invention of the electric battery by Volta made it
possible to study how matter behaves with electric current
-A gas discharge tube containing gas sealed under low
pressure could carry a current between two electrodes;
removing the gas from the tube did not stop the current;
this lead to the discovery of the cathode ray tube where a
current through the discharge tube came from the
cathode
-J. J. Thomson explained that the cathode rays could travel
in straight lines unaffected by gravity. They moved like light
waves at incredible fast speeds but could be deflected by
a magnet meaning they were charged.
-George Johnstone Stoney gave the negatively charged
particles the name electrons
-Thomson found that the charge of electrons is great
compared to its mass and that these particles existed in all
types of atoms. He created a model similar to plumpudding, where the negative electrons were embedded in
a pudding of positive protons.
-Wilhelm Roentgen discovered x-rays in 1896 which lead to
the discovery of nuclear radiation
-Ernest Rutherford experimented with alpha particles which
are positively charged ions with a mass of about 7300
electrons. They are emitted at high speeds from
radioactive particles. Their large mass and high speed
gives them great energy; have a +2 charge.
-Hans Geiger (Rutherford’s assistant) designed an
experiment that discovered the nucleus. Alpha particles
were beamed at a thin sheet of gold foil. Most (20,000)
passed straight through, but a few were deflected at a
large angle or ricocheted back towards the source. This
led to the conclusion that most of the atom was empty
except for a small, dense region 1/100 000 the size of the
atom.
-Later discoveries showed that the nucleus is positive
and contains protons which are 1836 the size of electrons
-The number of protons are related to the atom’s
chemical and physical properties; given an atomic
number to recognize the number of protons
-In 1932, James Chadwick observed neutral particles
radiating from metal bombarded by alpha particles;
called neutrons; their mass is 1839 times an electron; in
nucleus
-Spectroscopy, the study of light, helped describe the
movement of electrons in the atom
-A continuous spectrum of light is emitted from matter
that is heated; each element has its unique line spectrum
-Neils Bohr devised a mathematical model that
suggested electrons exist in energy levels around the
nucleus called principal energy levels. Energy levels
depend on distance from the nucleus. When energy
excites the electrons, they “jump” from one level to a
higher one. The change in energy is distinct or quantized.
The movement of electrons creates the line spectra
because excess energy is radiated as electromagnetic
energy as electrons fall back to their original energy level.
-Lowest energy state of an atom, called the ground
state, contains 6-7 levels
-Einstein’s theories suggested that light consisted of
photons
-Louise de Broglie suggested that if waves could
behave like particles then particles could behave like
waves (de Broglie’s hypothesis) ; all forms of matter act
like waves; wavelength depends on mass and speed
-Electrons are both waves and particles; known as
wave-particle duality
-Werner Heisenberg determined that it is impossible to
know both the exact location and the exact energy of an
atom at the same time; Heisenberg uncertainty principle
-Orbits were replaced by orbitals which are 3D
regions of an electron’s probable location; look like fuzzy
clouds without definite boundaries
1. The law of definite composition states that every
compound has a definite composition by mass. What does
that mean?
The ratios of the masses that make up the compound are
the same no matter how much of the compound is used.
2. Upon what evidence was Dalton’s model of the atom
based? Compare that evidence with the evidence the
Greeks used to support their model of the atom. Are these
models directly related?
Dalton’s model was based on experimental evidence
collected over decades; the Greek model was based on
philosophical ideas. These models are not related because
their conclusions are based on two different methods of
inquiry.
3. What inventions permitted the investigation of the nature
of atoms?
The electrical battery and the gas discharge tube.
4. What four characteristics did Thomson and others
determine about cathode rays that resulted in the
discovery of the electron?
1. cathode rays were tiny particles, 2. they were
negatively charged, 3. the same kind of particle was
emitted by all kinds of atoms, 4. their charge compared
to their mass was extremely large
5. How did Thomson’s discovery change the model of
the atom?
The atom was no longer viewed as a solid, indestructible
sphere, but rather as an assemblage of parts that could
be changed under certain conditions.
6. After considering the evidence of his assistant’s
experiment, how did Rutherford describe the atom and
on what evidence was his model based?
Since most of the alpha particles pass undeflected
through the gold, he concluded that most of the atom
was empty space. The small number of particles that
were greatly deflected must have struck a small, very
dense region.
7. Why was the discovery of the neutron difficult? Why was
it important for the development of the atomic model?
It is neutral so it could not be deflected by magnetic fields
or charged plates. Its discovery accounted for the extra
mass in the nucleus.
8. What does it mean to say that an electron’s energy within
an atom is quantized?
The levels in which electrons reside represent only certain
energies, no more and no less. When they change energy
levels they absorb and emit energy in only discrete amounts.
9. Describe the main difference between the standard
Bohr model of the electron and the quantum model.
In the Bohr model, electrons orbit the nucleus like a
planet at certain distances. In the quantum model,
electrons behave like matter waves that occupy a three
dimensional space surrounding the nucleus.
10. (true/false) One of the most revolutionary aspects of
Dalton’s model of the atom was that the atoms of each
element had masses that were different from the masses
of atoms of other elements.
True
-Every electron has its own energy-determined location
within the atom
-The location of the electron is given by 4 quantum
numbers
-Quantum numbers are solutions to the various wave
equations scientists use to describe the energy,
momentum, and probable location of an electron
-The principal quantum number identifies the principal
energy level which is the average or most likely distance
an electron is from the nucleus
-Also known as the energy-level number
-Positive integers beginning at 1
-Highest value is 7 in the ground state
-Can be higher in excited atoms
-All electrons in a given energy level have the same
principal quantum number
-Total electrons in a principal energy level equals 2n2
-Within the energy levels, electrons have a greater
probability of being found in certain sublevels
-Also called the sublevel quantum number
-Azimuth refers to a direction
-Azimuthal quantum numbers are solutions to angular
motion equations that describe these sublevels
-Sublevels are normally symmetric in shape
-Shapes are identified by s, p, d, and f
-For each value of n, there are n possible l (Azimuthal)
numbers
-l can have values from 0 to n-1
-Values of l are designated by letters s=0, p=1, d=2, and f=3
-Orientations of electrons are effected by magnetic
fields
-The magnetic quantum number describes the spatial
orientation of the orbitals within the atoms
-Also called the orbital quantum number
-Can be small negative and positive numbers
-On each sublevel there are 2l + 1 possible orbitals (m
values)
-m is represented by 0, +-1, +-2, +-3, …
-Each orbital can hold only two electrons
-These electrons have the opposite reactions to a
magnetic field; called a “spin” although their movement
is wavelike
-Two possible spin values +1/2 represented by an up
arrow and -1/2 represented by a down arrow
-Since the spin is always a different number, no two
quantum number addresses can be the same; known as
the Pauli exclusion principle
-Each principal energy level (besides the first) is divided
into sublevels which are assigned l values
-Sublevels are divided into orbitals assigned m values
-Orbitals are divided by their spin values ms
-The sublevels l are named s, p, d, and f with s=0, p=1,
d=2, and f=3
- The s sublevel
-Simpliest, only one orbital (2 electrons)
-Spherical shape
-Occurs when there is only one energy level
- The p sublevel
-Dumbbell-shaped, three dimensional on a x-y-z graph
-p orbitals labeled as px , py , and pz
-Two electrons can be in each of the 3 orbitals making
a total of 6 electrons
-Occurs when there are 2 or more energy levels
- The d sublevel
-5 orbitals making 10 electrons
-Occurs when there are 3 or more energy levels
- The f sublevel
-7 orbitals making 14 electrons
-Occurs when there are 4 or more energy levels
Allowed Electron Capacities of Energy Levels
principal energy level
allowed sublevels
orbitals in each level
electron capacity
total electrons
1
s
1
2
2
2
s
1
2
8
p
3
6
s
1
2
p
3
6
d
5
10
s
1
2
p
3
6
f
5
10
d
7
14
s
1
2
p
3
6
d
5
10
f
7
14
s
1
2
p
3
6
d
5
10
s
1
2
p
3
6
3
4
5
6
7
18
32
32
18
8
-Principal energy levels become larger in size as their
energy levels increase (further from nucleus)
-Sublevels can also be ranked by energy
-1s has the least energy then 2s and 2p
-Then it is 3s and 3p
-However, since electrons fill 8 electrons on a level and
then move up then next lowest is 4s not 3d
-Follows the diagonal rule (page 88)
-Periodic table is organized by electron configuration
-Each row corresponds to the highest principal energy
level
-Blocks of the table indicate which sublevel electrons
are added
-s will hold only 2; first two columns are the s block
-p will hold 6; last six columns are the p block
-d block is in the middle (will be filled after s orbitals)
-f block at the bottom (will be filled after d orbitals)
-Atomic numbers tell the numbers of electrons
-Aufbau principle: arrangement of an atom is
determined by adding electrons to an atom with a lower
atomic number; lower energy orbitals are filled first
-Electron configurations: arrangements of electrons
-Electron configurations are written with a coefficient
indicating the energy level, the letter of the sublevel, and
a subscript to tell how many electrons are in that sublevel
-Examples: H is 1s1 and He is 1s2
-Lithium is 1s22s1 and beryllium is 1s22s2
-Orbital notation has a horizontal line to represent the
orbital with the orbital name written above; an up arrow
is used for +1/2 and a down arrow for -1/2
-Instead of copying the entire electron configuration
for the element filling the previous energy level, its symbol
is put in a bracket
-Lithium could be written [He] 2s1
-In placing the electrons in orbital notation, all the
orbitals receive one electron before pairs are created;
known as Hund’s rule
-Raymond Damadian developed MRI (magnetic
resonance imaging) in the 1970s
-Based on magnetic fields created in atoms
-Places body in magnetic field which aligns hydrogen
nuclei in body
-Radio frequency makes hydrogen wobble like a
spinning top
-When the radio frequency is turned off, the hydrogen
relax and realign which releases energy
-The released energy can be interpreted as a
computer image
1. List the formal names and symbols for the four quantum
numbers, their alternate names (if applicable), and all
possible values in the ground state.
Principal quantum number (n); energy level number; n =
1, 2, 3, …, 7
Azimuthal quantum number (l); sub-level quantum
number; l = 1, 2, 3, … , n-1
Magnetic quantum number (m); orbital quantum
number; for each l number, the possible m numbers are 0
and +-1, +-2, +-3, …, +-l
Electron spin quantum number (ms), +-1/2
2. To which quantum numbers do the sublevel
designations s, p, d, and f correspond? Give the
maximum electron capacity of each.
The s sublevel corresponds to l = 0, p to l = 1, d to l = 2,
and f to l = 3. The s sublevel may contain up to 2
electrons, the p up to 6 electrons, the d up to 10
electrons, and the f up to 14 electrons.
3. What do the Aufbau principle and Hund’s rule tell us
about the electron structure of an atom in its ground
state?
The Aufbau principle states that the electron structure of
the ground state atom may be determined by the
addition of electrons to a smaller atom. Additional
electrons occupy increasing higher sublevels. Hund’s rule
states that, for sublevels with more than one orbital, the
lowest energy configuration occurs when the most
unpaired electrons with the same spin occupy the
orbitals of the sublevel. (All orbitals get one electron
before any get two)
-Every electron has its own energy-determined location
within the atom
-The location of the electron is given by 4 quantum
numbers
-Quantum numbers are solutions to the various wave
equations scientists use to describe the energy,
momentum, and probable location of an electron
4. (true or false) The electron structure of an atom is
determined by direct observation using x-rays.
False - by theoretical modeling and emission spectra
5. (true or false) Only an odd number of orbitals is possible
for any electron sublevel.
True
6. For each of the following atoms, (1) give the
abbreviated ground-state electron configuration, (2)
draw the abbreviated orbital notation, and (3) find the
set of four quantum numbers for the last electron added.
a. silicon (Z=14)
(1) [Ar] 3d8 4s2
3d
(2) Si: [Ne] ___ ___ ___ ___ ___
(3) n = 3, l = 1, m = 0, ms = +1/2
4s
____
6. For each of the following atoms, (1) give the
abbreviated ground-state electron configuration, (2)
draw the abbreviated orbital notation, and (3) find the
set of four quantum numbers for the last electron added.
b. nickel (Z=28)
(1)[Ar] 3d8 4s2
3d
(2) Ni: [Ar] ___ ___ ___ ___ ____
(3) n = 3, l = 2, m = 0, ms = -1/2
4s
___
7. Tell whether each of the following quantum numbers,
given in the order n, l, m, ms, is possible or impossible. If
impossible, tell why.
a. (0, 1, 0, +1/2) impossible, n must be greater of equal to 1
b. (4, 1, 0, -1/2)
possible
c. (2, 0, 1, +1/2)
impossible, m can only be o for l = 1
8. Moving electrical charges create magnetic fields. Why
is a hydrogen nucleus charged? Do you think electrons
also generate magnetic fields?
Hydrogen nuclei are charged because of the positive
charge of the protons that inhabit nuclei. Electrons also
generate magnetic fields because they “spin.”
-Scientists agree that the basic subatomic particles are
the proton, neutron, and the electron
-Atomic number equals the number of protons
-Represented by Z
-Size of positive charge determines the number of
electrons
-Number of electrons determines the chemical
properties
-Numbers of neutrons made the mass of different atoms
of the same element different
-Atoms of the same element are not all identical;
neutron number is different
-Isotope: same number of protons but different
numbers of neutrons
-Chemical properties of isotopes are the same
-Mass number is a way to identify the isotope; equal to
atomic number Z + neutrons (N)
-Isotopic notation indicates the specific isotope in a
chemical expression; written as a superscript mass
number over a subscript atomic number followed by the
elements abbreviation
-Mass is based on 1/12 the weight of a carbon-12 atom;
defined as an unified atomic mass unit (amu)
-Not all atoms have the same mass since isotopes with
more neutrons weigh more; use a weighted average of
all naturally occurring isotopes
-Example:
Calculate the average atomic mass of magnesium with
isotopes Mg-24 (23.985 u) at 78.99%, Mg-25 (24.986 u) at
10.00%, and Mg-26 (25.983 u) at 11.01%.
24.3042 u = 24.30 u
-Electrons are responsible for an atom’s chemical
behavior
-Electrons in the outermost energy levels are most likely
to be involved in chemical bonding; they give elements
their chemical properties; known as valence electrons
-Valence electrons have the highest principal quantum
number (Coefficient in front of the sublevel letter)
-Example:
Find the number of valence electrons for the following
elements.
a. Argon
Z=18
[Ne] 3s2 3p6
Valence electrons : 8
b. Nickel
Z=28
[Ar] 3d8 4s2
Valence electrons : 2
-Electron-dot notation is way to represent the number of
valence electrons
-Dots placed around symbol; no more than 8; only refer
to p and s sublevels
-Order that dots are placed
3 6
4
2
7
1
5 8
-The 3rd – 8th electrons (ones in the p sublevels) follow
Hund’s rule and do not pair up until all are in place
-Electrons are gained or lost creating ions
-Ions are charged particles
-Cations are positive ions with more protons than
electrons
-Anions are negative ions with more electrons than
protons
-A charge is represented by a negative or positive
superscript
-Example:
Write the electron dot notation for the following
elements.
a. sulfur
Z=16
[Ne] 3s2 3p4
b. boron
Z=5
[He] 2s2 2p1
c. chlorine anion (Cl-)
Z=17
[Ne] 3s2 3p5
1. Identify the symbol or notation that is used to represent
the following information.
a. atomic number
Z
b. mass number
A
c. number of neutrons
N
d. identification of different isotopes
isotopic notation
e. configuration of valence electrons
electron-dot notation
2. Compare and contrast an element’s mass number
and its atomic mass.
Mass number is a whole number equal to the sum of the
protons and neutrons in a certain isotope. Atomic mass is
a decimal that is equal to weighted average of all
isotopes of a given element.
3. Determine the number of protons, neutrons, and
electrons in the neutral atoms of the following elements.
a. silicon-28
b. calcium-44
Z = 14
Z = 20
protons = 14
protons = 20
neutrons = 14
neutrons = 24
electrons = 14
electrons = 20
4. How is the current atomic mass in an atom defined?
1/12 the mass of the carbon-12 isotope
5. Which electrons in an atom are its valence electrons?
Why are these electrons significant?
The ones that occupy the highest energy level. They are
responsible for almost all of the atoms chemical
properties.
6. Write the abbreviated (noble gas) electron
configuration for the elements listed below. How many
valence electrons does each have?
a. carbon C (Z = 6)
[He] 2s2 2p2
4 valence eb. fluorine C (Z = 9)
[He] 2s2 2p5
7 valence e-
6. Write the abbreviated (noble gas) electron
configuration for the elements listed below. How many
valence electrons does each have?
c. potassium C (Z = 19)
[Ar] 4s1
1 valence ed. krypton C (Z = 36)
[Ar] 3d10 4s24p6
8 valence e-
7. Explain the difference between a positive ion and a
negative ion?
A positive ion has fewer electrons than protons and a
negative ion has more electrons than protons.
8. (true or false) Isotopes of a given element are
chemically identical but may have different physical or
nuclear properties (ex radioactivity).
True
9. Write the electron dot notations for all the elements in
the following columns of the periodic table.
a. column 2
b. column 15
Be
N
Mg
P
Ca
As
Sr
Sb
Ba
Bi
Ra
What generalization seems to be true about the electrondot notation for elements in a column of the periodic
table?
It is the same