Physical science unit 2
Greek thinkers, such as Democritus. He
called nature’s basic particle an atom,
based on the Greek word meaning
“indivisible.” The particle theory of matter
was supported as early as 400 B.C. by
certain
Aristotle was part of the generation that
succeeded Democritus. His ideas had a
lasting impact on Western civilization,
and he did not believe in atoms.
He thought that all matter was
continuous, and his opinion was
accepted for nearly 2000 years.
Neither the view of Aristotle nor that of
Democritus was supported by
experimental evidence, so each
remained speculation until the
eighteenth century. Then scientists
began to gather evidence favoring the
atomic theory of matter
Aided by improved balances,
investigators began to accurately
measure the masses of the elements and
compounds they were studying. This led
to the discovery of several basic laws:
law of conservation of mass, which
states that mass is neither created
nor destroyed during ordinary
chemical reactions or physical
changes
law of definite proportions
a chemical compound contains the same
elements in exactly the same proportions
by mass regardless of the size of the
sample or source of the compound
For example, sodium chloride, also
known as ordinary table salt, always
consists of 39.34% by mass of the
element sodium, Na, and 60.66% by
mass of the element chlorine, Cl
law of multiple proportions:
If two or more different compounds are
composed of the same two elements,
then the ratio of the masses of the
second element combined with a certain
mass of the first element is always a
ratio of small whole numbers.
Ex: the elements carbon and oxygen
form two compounds, carbon dioxide and carbon monoxide.
In carbon dioxide, 2.66 g of oxygen combine with 1.00 g of
carbon.
In carbon monoxide, 1.33 g of oxygen combine with 1.00 g of
carbon
The ratio of the masses of oxygen in these two compounds is
2.66 to 1.33, or 2 to 1
In 1808, an English schoolteacher
named John Dalton proposed an
explanation for the law of conservation
of mass, the law of definite proportions,
and the law of multiple proportions
Dalton’s Atomic Theory
• 1. All matter is composed of extremely
small particles called atoms.
Dalton’s Atomic Theory
• 2. Atoms of a given element are identical
in size, mass, and other properties; atoms
of different elements differ in size, mass,
and other properties.
Dalton’s Atomic Theory
• 3. Atoms cannot be subdivided, created,
or destroyed.
anyone have a problem with this one?
Dalton’s Atomic Theory
• 4. Atoms of different elements combine in
simple whole-number ratios to form
chemical compounds.
Dalton’s Atomic Theory
• 5. In chemical reactions, atoms are
combined, separated, or rearranged.
Discovery of the electron
• The first discovery of a subatomic particle
resulted from investigations into the
relationship between electricity and matter.
In the late 1800s, many experiments were
performed in which electric current was
passed through various gases at low
pressures. (Gases at atmospheric
pressure don’t conduct electricity well.)
Cathode-ray tube experiment
• Investigators noticed that when current
was passed through a cathode-ray tube,
the surface of the tube directly opposite
the cathode glowed.
Cathode ray
• They hypothesized that the glow was
caused by a stream of particles
Cathode rays were deflected by a magnetic field in
the same manner as a wire carrying electric
current, which was known to have a negative
charge
So the rays were deflected away from a
negatively charged object
the particles that compose cathode
rays are negatively charged….
Like charges repel each other.
(1897) English physicist Joseph John
Thomson measured the ratio of the charge of
cathode-ray particles to their mass. He found
that this ratio was always the same, regardless
of the metal used to make the cathode or the
nature of the gas inside the cathode-ray tube.
Thomson concluded that all cathode rays are
composed of identical negatively charged
particles, which were named “electrons”
This led to other questions.
• 1. Because atoms are electrically neutral,
they must contain a positive charge to
balance the negative electrons.
• 2. Because electrons have so much less
mass than atoms, atoms must contain
other particles that account for most of
their mass.
Thomson proposed the plum pudding model .
He believed that the negative electrons were
spread evenly throughout the positive charge of
the rest of the atom (but do not contribute much
to the overall mass)
new experiments would disprove this model
1909 Robert Millikan
• Oil drop experiment- Millikan was able to
calculate the charge on an electron by
giving a drop of oil a negative charge and
then having it fall towards a negatively
charged plate.
• http://www.learnerstv.com/animation/anim
ation.php?ani=186&cat=chemistry
(1911) Ernest Rutherford bombarded a
thin piece of gold foil with fast-moving
alpha particles, which are positively
charged particles with about four times
the mass of a hydrogen atom.
Rutherford- gold foil
• The particles were expected to
pass through the foil with little to
no deflection.
but…………..
Rutherford was quoted,
“it was as if you had fired a 15-inch
artillery shell at a piece of tissue paper
and it came back and hit you.”
Rutherford- gold foil
This led him to believe the following:
Whatever is repelling the alpha particles is
positively charged but it does not take up a
lot of space, because most alpha particles
get through.
Alpha
particle
source
So:
all atomic nuclei are made of two kinds of
particles, protons and neutrons.
A proton has a positive charge equal in
magnitude to the negative charge of an
electron.
Atoms are electrically neutral because
they contain equal numbers of protons
and electrons.
(A neutron is electrically neutral)
A proton has a mass of 1.673 × 10−27
kg, which is 1836 times greater than
the mass of an electron
Like charges normally repel each other, but
when protons are very close together they are
actually attracted to one another.
As many as 83 protons can exist close
together to help form a stable nucleus
In larger atoms the nucleus is unstable
A similar attraction exists when neutrons
are very close to each other or when
protons and neutrons are very close
together.
short-range proton-neutron, protonproton, and neutron-neutron forces
hold the nuclear particles together
and are referred to as nuclear
forces.
We sometimes refer to the region
occupied by the electrons as an
electron cloud —a cloud of negative
charge.
The atomic radius is the distance
from the center of the nucleus to the
outer portion of this electron cloud.
Because atomic radii are so small, they
are expressed using a unit that is more
convenient for the sizes of atoms. This
unit is the picometer.
The abbreviation for the picometer is pm
(1 pm = 10−12 m = 10−10 cm).
Atomic radii range from about 40 to 270
pm.
Nuclei also have incredibly high
densities,
about
200 000 000metric tons/cm3.
Atoms of different elements have
different numbers of protons.
Atoms of the same element all have
the same number of protons
The atomic number of an element is
the number of protons of each atom
of that element
Elements are placed on the periodic table
in order of increasing atomic number.
At the top left of the table is hydrogen, H,
which has atomic number 1.
Next in order is helium, He, which has
two protons.
Lithium, Li, has three protons
All hydrogen atoms have only one
proton.
However, like many naturally
occurring elements, hydrogen atoms
can have different numbers of
neutrons, but still be of the same
element.
Isotopes are atoms of the same
element that have different masses.
The isotopes of a particular element
all have the same number of protons
and electrons but different numbers
of neutrons.
Example:
Three types of hydrogen atoms are
known. The most common type
of hydrogen is sometimes called protium.
It accounts for 99.9885% of the hydrogen
atoms found on Earth. The nucleus of a
protium atom consists of one proton only,
and it has one electron moving about it.
Another is called deuterium, which
accounts for 0.0115% of Earth’s
hydrogen atoms. Each deuterium
atom has a nucleus with one proton
and one neutron
The third form of hydrogen is known
as tritium, which is radioactive. It
exists in very small amounts in
nature, but it can be prepared
artificially. Each tritium atom has one
proton, two neutrons, and one
electron.
In all three isotopes of hydrogen, the
positive charge of the single proton
is balanced by the negative charge
of the electron.
So, once again, what is the
difference between the three?
mass
• Although isotopes have different
masses, they do not differ
significantly in their chemical
behavior.
The mass number is the total
number of protons and neutrons that
make up the nucleus of an isotope
The three isotopes of hydrogen
described earlier have mass
numbers 1, 2, and 3
Identifying an isotope requires
knowing both the name or atomic
number of the element and the
mass of the isotope
Designating Isotopes
There are two methods for specifying
isotopes. In the first method (called
hyphen notation) the mass
number is written with a hyphen after
the name of the element. Tritium,
for example, is written as hydrogen-3.
The second method shows the
composition of a nucleus using the
isotope’s nuclear symbol. For
example, uranium-235 is written as
235
U92 U
The superscript indicates the mass
number (protons + neutrons) and
the subscript indicates the atomic
number (number of protons).
From the nuclear symbol the
number of neutrons is found by
subtracting the atomic number
from the mass number.
235
U92 U
235 (protons + neutrons) − 92 (protons) = 143 neutrons
How many protons, neutrons and electrons
make up an atom of bromine-80?
answer
• Protons-35 (the atomic number is 35,
which indicates how many protons per
atom)
• Electrons-35 (protons must be = to
electrons or else the atom cannot be
electrically neutral)
• Neutrons-45 (mass number 80- atomic
number 35)
Write the nuclear symbol for carbon-13
answer
13
6
C
Write the hyphen notation for the isotope
with 15 electrons and 15 neutrons
answer
15 electrons means 15 protons, so it has to
be phosphorus. 15 neutrons added to the
15 protons gives it a mass number of 30
phosphorus-30
Nuclide is a general term for a
specific isotope of an element
we have discussed 3 nuclides of
Hydrogen
Relative Atomic Masses
Masses of atoms expressed in grams
are very small. An atom of oxygen-16, for
example, has a mass of
2.656 × 10–23 g.
For most chemical calculations it is more
convenient to use relative atomic
masses.
one atom has been arbitrarily chosen
as the standard and assigned a mass
value. The masses of all other atoms
are expressed in relation to this defined
standard. The standard used by
scientists to compare units of atomic
mass is the carbon-12 atom. It has
been arbitrarily assigned a mass of
exactly 12 atomic mass units, or 12
amu.
One atomic mass unit, or 1 amu,
is exactly 1/12 the mass of a
carbon-12 atom
The atomic mass of any other atom is
determined by comparing it with the
mass of the carbon-12 atom.
The hydrogen-1 atom has an atomic
mass of about 1/12 that of the
carbon-12 atom.
So:
a Hydrogen-1 atom has a mass of 1amu
• How much does 1 atom of oxygen weigh?
16amu
To be exact:
The mass of the electron is
0.000 5486 amu,
The mass of the proton is
1.007 276 amu,
The mass of the neutron is
1.008 665 amu.
Average Atomic Masses of
Elements
Average atomic mass is the
weighted average of the atomic
masses of the naturally occurring
isotopes of an element
The percentage of each isotope in the
naturally occurring element on Earth is
nearly always the same, no matter
where the element is found. The
percentage at which each of an
element’s isotopes occurs in nature is
taken into account when calculating the
element’s average atomic mass
The average atomic mass of an element
depends on both the mass and the relative
abundance of each of the element’s isotopes.
For example, naturally occurring
copper consists of 69.15% copper63 (which has an atomic mass of
62.929 601 amu) and 30.85%
copper-65 (which has an atomic
mass of 64.927 794 amu)
The average atomic mass of the
element can be calculated by
multiplying the atomic mass of
each isotope by its relative
abundance (expressed in decimal
form, not %) and adding the
results.
% ab C-63
mass C-63
% ab C-65
mass C-65
0.6915 × 62.929 601 amu + 0.3085 × 64.927 794 amu =
The calculated average atomic mass of naturally
occurring copper is 63.55 amu.
(We usually round avg atomic mass to two decimal
places)
The relative atomic mass scale
makes it possible to know how many
atoms of an element are present in a
sample of the element with a
measurable mass
Three very important concepts—the
mole, Avogadro’s number, and
molar mass make it possible to
relate mass in grams to number of
atoms.
The mole is the SI unit for amount of
substance. A mole (abbreviated
mol) is the amount of a substance
that contains as many particles as
there are atoms in exactly 12 g of
carbon-12
This is huge!
We now have a ratio for number of particle
to a given mass. That means we will be
able to convert between the two.
The number of particles in a mole has
been experimentally determined
in a number of ways. The best modern
value is 6.02 ×1023
This means that exactly 12 g sample
of carbon-12 contains:
6.022 141 79 ×1023 carbon-12 atoms
Amedeo Avogadro 1776-1856
Its hard to be this
good looking…….
Amedeo Avogadro 1776-1856
• Avogadro was a lawyer, physicist and
mathematician who (among other things)
made many contributions to molecular
theory.
Avogadro’s number—
6.022 141 79 × 1023—
is the number of particles in exactly
one mole of a pure substance. For
most purposes, Avogadro’s number
is rounded to 6.02 × 1023.
Avogadro’s number can be used to:
1. find the number of atoms of an
element from the amount in moles
2. find the amount of an element in
moles from the number of atoms
3. Convert between amount (in
moles) and mass in grams
Lets make a “mole map” to help
visualize this
“All roads go through
the mole”
How much does 6.951 mol of Fe
weigh?
6.951molFe x ______________ =
gFe
How much does 6.951 mol of Fe
weigh?
6.951molFe x
55.85g
______________
=
1 mol Fe
388.2 gFe
How many moles of Zn are there in
23.9g?
23.9gZn x ______________ =
molZn
How many moles of Zn are there in
23.9g?
1 mol Zn
23.9gZn x ______________ =
65.39g
0.366 molZn
How many atoms are there in
1.999 mol Li?
1.999molLi x ______________ =
atoms
How many atoms are there in
1.999 mol Li?
6.02x1023 atoms
1.999molLi x ______________
1 molLi
= 1.2x1024 atoms
The arrangement of electrons in
atoms
The Rutherford model of the atom was an
improvement over previous models, but it was
incomplete. It did not explain how the atom’s
negatively charged electrons are distributed in
the space surrounding its positively charged
nucleus. After all, it was well known that
oppositely charged particles attract each
other. So what prevented the negative
electrons from being drawn into the positive
nucleus?
In the early twentieth century, a new atomic
model evolved as a result of investigations
into the absorption and emission of light by
matter. The studies revealed a relationship
between light and an atom’s electrons. This
new understanding led directly to a
revolutionary view of the nature of energy,
matter, and atomic structure
Before 1900, scientists thought light behaved
solely as a wave. This belief changed when it
was later discovered that light also has particlelike characteristics. Still, many of light’s
properties can be described in terms of waves.
A quick review of these wavelike properties will
help you understand the basic theory of light as
it existed at the beginning of the twentieth
century.
The Wave Description of Light
Visible light is a kind of electromagnetic
radiation, which is a form of energy that
exhibits wavelike behavior as it travels
through space. Other kinds of
electromagnetic radiation include X rays,
ultraviolet and infrared light, microwaves,
and radio waves. Together, all the forms of
electromagnetic radiation form the
electromagnetic spectrum.
All forms of electromagnetic radiation
move at a constant speed of
3.00 × 108 meters per second (m/s)
through a vacuum and at slightly
slower speeds through matter.
Because air is mostly empty space, the
value of 3.00 × 108 m/s is also light’s
approximate speed through air.
The significant feature of wave motion is its
repetitive nature, which can be characterized
by the measurable properties of wavelength
and frequency
Wavelength (λ) is the distance between
corresponding points on adjacent
waves. The unit for wavelength is a
distance unit. Depending on the type of
electromagnetic radiation, it may be
expressed in meters, centimeters, or
nanometers
Frequency (ν) is the number of waves that
pass a given point in a specific time, usually
one second.
Frequency is expressed in waves/second
One wave/second is called a hertz (Hz),
named for Heinrich Hertz, who was a
pioneer in the study of electromagnetic
radiation.
l
v
Frequency and wavelength are
mathematically related to each other.
For electromagnetic radiation, this
relationship is written as follows.
c = λν
C=(constant)speed of light
λ= wavelength
V= frequency
In the equation,
c is the speed of light (in m/s),
λ is the wavelength of the electromagnetic
wave (in m),
ν is the frequency of the electromagnetic
wave (in s−1)
Because c is the same for all
electromagnetic radiation, the product λν
is a constant. Consequently, we know that
λ is inversely proportional to ν. In other
words, as the wavelength of light
decreases, its frequency increases, and vice
versa.
In the early 1900s, scientists conducted
two experiments involving interactions
of light and matter that could not be
explained by the wave theory of light. One
experiment involved a phenomenon
known as the photoelectric effect
The photoelectric effect refers to the emission
of electrons from a metal when light shines on
the metal
The mystery of the photoelectric effect involved the
frequency of the light striking the metal. For a given
metal, no electrons were emitted if the light’s
frequency was below a certain minimum—
regardless of the light’s intensity. Light was known to
be a form of energy, capable of
knocking loose an electron from a metal. But the
wave theory of light predicted that light of any
frequency could supply enough energy to eject an
electron. Scientists couldn’t explain why the light
had to be of a minimum frequency in order for the
photoelectric effect to occur.
The Particle Description of Light
The explanation of the photoelectric effect
dates back to 1900, when German physicist
Max Planck was studying the emission of light
by hot objects. He proposed that a hot object
does not emit electromagnetic energy
continuously, as would be expected if the
energy emitted were in the form of waves.
Instead, Planck suggested that the object
emits energy in small, specific packets called
quanta.
A quantum of energy is the minimum
quantity of energy that can be lost or
gained by an atom.
Planck proposed the following relationship
between a quantum of energy and the
frequency of radiation.
E = hν
In the equation,
E =energy, in joules, of a quantum of
radiation
v =frequency, in s−1, of the radiation
emitted
h = a fundamental physical constant now
known as Planck’s constant
h = 6.626 × 10−34 J• s
In 1905, Albert Einstein expanded on
Planck’s theory by introducing
the radical idea that electromagnetic
radiation has a dual wave-particle
nature. While light exhibits many wavelike
properties, it can also be thought of as a
stream of particles. Each particle of light
carries a quantum of energy. Einstein called
these particles photons
A photon is a particle of electromagnetic
radiation having zero mass and carrying a
quantum of energy
The energy of a particular photon depends
on the frequency of the radiation.
Ephoton = hν
Einstein explained the photoelectric effect by
proposing that electromagnetic
radiation is absorbed by matter only in whole
numbers of photons. In order for an electron to
be ejected from a metal surface, the electron
must be struck by a single photon possessing at
least the minimum energy required to knock
the electron loose. According to the
equation Ephoton = hν, this minimum energy
corresponds to a minimum frequency
If a photon’s frequency is below the
minimum, then the electron remains
bound to the metal surface. Electrons in
different metals are bound more or less
tightly, so different metals require different
minimum frequencies to exhibit the
photoelectric effect.
When current is passed through a gas at
low pressure, the potential energy of the
gas atoms increases.
The lowest energy state of an atom is its
ground state.
A state in which an atom has a higher
potential energy than it has in its ground
state is an excited state
There are many possible excited
states, each with a unique energy, but only
one ground state energy for atoms of a
given element. When an excited atom
returns to its ground state or a lower
energy excited state, it gives off the energy
it gained in the form of electromagnetic
radiation
When investigators passed electric current
through a vacuum tube containing
hydrogen gas at low pressure, they
observed the emission of a characteristic
pinkish glow. When a narrow beam of the
emitted light was shined through a prism, it
was separated into four specific colors of
the visible spectrum.
The four bands of light were part of what is known as
hydrogen’s emission-line spectrum
Classical theory predicted that the
hydrogen atoms would be excited by
whatever amount of energy was added to
them. Scientists had thus expected to
observe the emission of a continuous
range of frequencies of electromagnetic
radiation, that is, a continuous spectrum
Why had the hydrogen atoms given off
only specific frequencies of light? Attempts
to explain this observation led to an
entirely new atomic theory called
quantum theory.
Whenever an excited hydrogen atom falls
to its ground state or to a lower-energy
excited state, it emits a photon of
radiation.
The energy of this photon (Ephoton = hν) is
equal to the difference in energy between
the atom’s initial state and its final state
The puzzle of the hydrogen-atom spectrum
was solved in 1913 by the Danish physicist
Niels Bohr. He proposed a hydrogen-atom
model that linked the atom’s electron to
photon emission. According to the model,
the electron can circle the nucleus only in
allowed paths, or orbits. When the
electron is in one of these orbits, the atom
has a definite, fixed energy
The electron—and therefore the hydrogen
atom—is in its lowest energy state when it
is in the orbit closest to the nucleus. This
orbit is separated from the nucleus by a
large empty space where the electron
cannot exist. The energy of the electron is
higher when the electron is in orbits that
are successively farther from the nucleus.
The electron orbits, or atomic energy
levels, in Bohr’s model can be compared to
the rungs of a ladder. When you are
standing on a ladder, your feet are on one
rung or another. The amount of potential
energy that you possess corresponds to
standing on the first rung, the second
rung, and so forth
How does Bohr’s model of the hydrogen atom
explain the observed spectral lines? While in a
given orbit, the electron is neither gaining nor
losing energy. It can, however, move to a higherenergy orbit by gaining an amount of energy
equal to the difference in energy between the
higher-energy orbit and the initial lower-energy
orbit. When a hydrogen atom is in an excited
state, its electron is in one of the higher-energy
orbits.
When the electron falls to a lower energy
level, a photon is emitted, and the process
is called emission. The photon’s energy is
equal to the energy difference between the
initial higher energy level and the final
lower energy level.
Energy must be added to an atom in order
to move an electron from a lower energy
level to a higher energy level.
This process is called absorption
Absorption and emission of radiation
in Bohr’s model of the hydrogen atom are
illustrated. The energy of each absorbed or
emitted photon corresponds to a particular
frequency of emitted radiation,
Ephoton = hν.
Based on the different wavelengths of the
hydrogen emission-line spectrum, Bohr
calculated the allowed energy levels for the
hydrogen atom. He then related the
possible energy-level changes to the lines
in the hydrogen emission-line spectrum
The Quantum Model
of the Atom
To the scientists of the early twentieth century,
Bohr’s model of the hydrogen atom
contradicted common sense. Why did
hydrogen’s electron exist around the nucleus
only in certain allowed orbits with
definite energies? Why couldn’t the electron
exist in a limitless number of orbits with slightly
different energies? To explain why atomic
energy states are quantized, scientists had to
change the way they viewed the nature of the
electron
Electrons as Waves
The investigations into the photoelectric
effect and hydrogen’s emission-line
spectrum revealed that light could behave
as both a wave and a particle. Could
electrons have a dual wave-particle nature
as well?
In 1924, the French scientist Louis de
Broglie asked himself this very
question. And the answer that he proposed
led to a revolution in our basic
understanding of matter. De Broglie
pointed out that in many ways the
behavior of electrons in Bohr’s quantized
orbits was similar to the known behavior of
waves
scientists at the time knew that any wave
confined to a space can have only certain
frequencies. De Broglie suggested that
electrons be considered waves confined to
the space around an atomic nucleus
It followed that the electron waves could
exist only at specific frequencies. And
according to the relationship E = hν, these
frequencies corresponded to specific
energies—the quantized energies of Bohr’s
orbits.
Other aspects of de Broglie’s hypothesis
that electrons have wavelike properties
were soon confirmed by experiments.
Investigators demonstrated that electrons,
like light waves, can be bent, or diffracted.
Diffraction refers to the bending of a wave
as it passes by the edge of an object or
through a small opening. Diffraction
experiments and other investigations also
showed that electron beams, like waves,
can interfere with each other
Interference occurs when waves overlap This
overlapping results in a reduction of energy in
some areas and an increase of energy in others.
The idea of electrons having a dual waveparticle nature troubled scientists.
If electrons are both particles and waves,
then where are they in the atom? To
answer this question, it is important to
consider a proposal first made in 1927 by
the German theoretical physicist Werner
Heisenberg.
Heisenberg’s idea involved the detection of
electrons. Electrons are detected by their
interaction with photons. Because photons
have about the same energy as electrons,
any attempt to locate a specific electron
with a photon knocks the electron off its
course. As a result, there is always a basic
uncertainty in trying to locate an electron
(or any other particle)
The Heisenberg uncertainty principle
states that it is impossible to determine
simultaneously both the position and
velocity of an electron or any other
particle. Although it was difficult for
scientists to accept this fact at the time, it
has proven to be one of the fundamental
principles of our present understanding of
light and matter.
The Schrödinger Wave Equation
In 1926, the Austrian physicist Erwin
Schrödinger used the hypothesis
that electrons have a dual wave-particle
nature to develop an equation
that treated electrons in atoms as waves.
John Q. Public
Don’t bother writing this one down………..
Unlike Bohr’s theory, which assumed
quantization as a fact, quantization of
electron energies was a natural outcome of
Schrödinger’s equation. Only waves of
specific energies, and therefore
frequencies, provided solutions to the
equation.
Together with the Heisenberg uncertainty
principle, the Schrödinger wave equation
laid the foundation for modern quantum
theory. Quantum theory describes
mathematically the wave properties
of electrons and other very small particles.
Solutions to the Schrödinger wave
equation are known as wave functions.
Based on the Heisenberg uncertainty
principle, the early developers
of quantum theory determined that wave
functions give only the probability of
finding an electron at a given place around
the nucleus.
Thus, electrons do not travel around the
nucleus in neat orbits, as Bohr
had postulated. Instead, they exist in
certain regions called orbitals
An orbital is a three-dimensional region
around the nucleus that indicates
the probable location of an electron.
In the Bohr atomic model, electrons of
increasing energy occupy orbits
farther and farther from the nucleus.
According to the Schrödinger equation,
electrons in atomic orbitals also have
quantized energies. An electron’s energy
level is not the only characteristic of an
orbital that is indicated by solving the
Schrödinger equation.
In order to completely describe orbitals,
scientists use quantum numbers.
Quantum numbers specify the
properties of atomic orbitals and the
properties of electrons in orbitals
The first three quantum numbers result
from solutions to the Schrödinger
equation. They indicate the main energy
level, the shape, and the orientation of an
orbital.
The fourth, the spin quantum number,
describes a fundamental state of the
electron that occupies the orbital
The principal quantum number (n)
indicates the main energy level occupied
by the electron
Values of n are positive integers only—1, 2,
3, and so on
(it’s a whole number)
As n increases, the electron’s energy and its
average distance from the nucleus increase
Except at the first main energy level,
orbitals of different shapes— known as
sublevels—exist for a given value of n.
The angular momentum quantum number
indicates the shape of the orbital.
For a specific main energy level, the
number of orbital shapes possible is equal
to n. (this is usually referred to as
“sublevel” )
s orbitals are spherical (contains 1 orbital)
p orbitals have dumbbell shapes (contains
3 orbitals)
d orbitals are more complex (contains 5
orbitals)
f orbital shapes are even more complex
(contains 7 orbitals)
• An orbital can hold up to 2 electrons
In the first energy level:
n=1
there is only one sublevel possible:
an s orbital.
the second energy level:
n=2
has two sublevels:
the s and p orbitals
The third energy level:
n=3
has three sublevels:
the s, p, and d orbitals
The fourth energy level:
n=4
has four sublevels:
the s, p, d, and f orbitals
In an nth main energy level,
there are n sublevels.
Each atomic orbital is designated by the
principal quantum number followed by the
letter of the sublevel.
For example:
the 1s sublevel is the s orbital in the first
main energy level
the 2p sublevel is the set of three p orbitals
in the second main energy level
a 4d orbital is part of the d sublevel in the
fourth main energy level
two dimensional illustration of something
that is supposed to be 3 dimensional………
There are three lobes of a p orbital
They extend along the x, y, or z axis of a
three dimensional coordinate system
There are therefore three p orbitals in
each p sublevel, which are designated as
px , py , and pz orbitals. The three p
orbitals occupy different regions of space
There are five different d orbitals in each
d sublevel.
There are seven different f orbitals in
each f sublevel.
Spin Quantum Number
has only two possible values
+½
-½
-This indicates the two fundamental spin
states of an electron in an orbital.
Spin Quantum Number
Think of an electron like the Earth spinning
on an axis. The electron exists in one of
two possible spin states, which creates a
magnetic field. To account for the magnetic
properties of the electron, theoreticians of
the early twentieth century created the
spin quantum number.
The electrons available to be lost,
gained, or shared in the formation of
chemical compounds are referred to as
valence electrons.
Valence electrons are often located in
incompletely filled main-energy levels
which are the outer most levels.
Electron configuration
The arrangement of electrons in an atom
Like all systems in nature, electrons in
atoms tend to assume arrangements that
have the lowest possible energies.
The lowest energy arrangement of the
electrons for each element is called the
element’s ground-state electron
configuration
Rules
Governing
Electron
Configurations
Aufbau principle
an electron occupies the lowest-energy
orbital that can receive it.
Pauli exclusion principle
no two electrons in the same atom can
have the same set of four quantum
numbers. The principal, angular
momentum, and magnetic quantum
numbers specify the energy, shape, and
orientation of an orbital.
Hund’s rule
orbitals of equal energy are each occupied
by one electron before any orbital is
occupied by a second electron, and all
electrons in singly occupied orbitals must
have the same spin state.
Representing
Electron
Configurations
Orbital Notation
In orbital notation, an unoccupied orbital is
represented by a line
with the orbital’s
name written underneath the line.
An orbital containing one electron is
represented as
.
An orbital containing two electrons
is represented as showing the electrons
paired and with opposite spin states
H___
1s
He___
1s
Electron-Configuration Notation
Electron-configuration notation
eliminates the lines and arrows of
orbital notation. Instead, the number of
electrons in a sublevel is shown by adding
a superscript to the sublevel designation
The hydrogen configuration
is represented by
1s1
The superscript indicates that one
electron is present in hydrogen’s 1s
orbital.
The helium configuration is represented by
1s2
Here the superscript indicates that there
are two electrons in helium’s 1s orbital.
2 elec
6 elec
10 elec
max
max
max
14 elec
max
Name
Sym
P+
N
E-
Calcium
Ca
20
20
20
Argon
Ar
18
22
18
Potassium
K
19
20
19
Boron
B
5
6
5
Silicon
Si
14
14
14
Gold
Au
79
118
79
Sodium
Na
11
12
11
Carbon
C
6
6
6
Uranium
U
92
146
92
Write electron configuration
notation and orbital notation for :
Li
Be
B
C
N
Ne
Na
1s
2s 2p
3s 3p 3d
4s 4p 4d 4f
5s 5p 5d 5f
6s 6p 6d
7s 7p
We will put all the rules together to
determine electron configurations of
different elements