Download Chapter 11

Chapter 11
Modern Atomic Theory
11.1 Rutherford’s Atom

The nuclear atom (atom with a
nucleus) resulted from Ernest
Rutherford’s Gold Foil
Experiment.
 Rutherford and his colleagues
were able to show that an atom
had a nucleus made of protons
and neutrons, and that the
atom was mostly empty space.
 Rutherford
was not able to determine what
was going on with the electrons.
 Rutherford suggested that the electron
were rotating around the nucleus like
planets revolving around the sun.
 The
problem with Rutherford’s thoughts on
electrons was that he couldn’t explain why
the positively charged nucleus did not
attract the negatively charged electrons.
11.2 Electromagnetic Radiation
 Electromagnetic
radiation is energy that is
transmitted from one location to another
by light.
 There are many different types of
electromagnetic radiation, all of which are
contained on the electromagnetic
spectrum.

The different types of electromagnetic radiation
are different in two ways: wavelength, frequency.
 The wavelength (λ) is the distance between two
consecutive peaks.
 The frequency (υ) indicates how many peaks
pass a certain point per given period of time.
 Waves also have speed; the speed of a wave
indicates how fast it travels.

Electromagnetic radiation can be thought of as a
wave that carries energy through space.
 Electromagnetic radiation doesn’t always
behave as a wave, sometimes it behaves as a
stream of particles.
 A beam of light can be thought of as a stream of
tiny packets of energy called photons.
 Different wavelengths of electromagnetic
radiation carry different amounts of energy.
11.3 Emission of Energy by
Atoms

When atoms receive energy from some source,
they become excited and release this energy by
emitting light.
 The energy emitted is carried away by a photon.
 A photon corresponds to the exact energy
change experienced by the atom.
 High energy photons correspond to short
wavelength light and low energy photons
correspond to long wavelength light.

Teacher's Domain
Fireworks
11.4 The Energy Levels of
Hydrogen
 An
excited atom can release some or all of
its excess energy by emitting a photon and
thus move to a lower energy state.
 A photon is a “particle” of electromagnetic
radiation.
 The lowest possible energy state of an
atom is called the ground state.

We can learn a great deal about the energy
states of any atom by observing the photons it
emits.
 Scientists are able to learn because the different
wavelengths of light carry different amounts of
energy per photon.
 The energy contained in a photon corresponds
to the change in energy that the atom
experiences going from an excited state to a
ground state.

The simplest atom to study is hydrogen,
because it contains only one proton and one
electron.
 When scientists studied hydrogen they found
that hydrogen only emits certain colors of visible
light.
 Because only certain photons are emitted, we
know that only certain energy changes are
occurring.
 Meaning hydrogen atoms have discrete energy
levels.
 Hydrogen
atoms always emit photons of
the same discrete colors.
 This means that hydrogen atoms have the
same set of discrete energy levels.
 The energy levels are quantized, that only
certain values are allowed.
 Scientists were surprised by the
quantization of the energy levels.
Hydrogen Line Emission
11.5 The Bohr Model of the Atom

Niels Bohr constructed a model of the hydrogen
atom with quantized energy levels the agreed
with the hydrogen emission results.
 Bohr’s model had an electron moving in circular
orbits corresponding to the allowed energy
levels.
 He suggested the electron could jump to a
different level by absorbing or emitting a photon
of light with the correct energy content.
Bohr’s model seemed
like it explained
everything.
 The electrons don’t
fall into the nucleus
because they have a
specific amount of
energy.
 It explained the line
emission spectrum of
hydrogen.

 Bohr’s
model was unfortunately incorrect.
 Bohr’s calculations only worked for the
hydrogen atom.
 His calculations could not explain the line
emission spectrum of helium.
 The current model of the atom is not the
Bohr model.
 In the current model, electrons do not orbit
the nucleus like planets orbiting the sun.

Teacher’s Domain Video
11.6 The Wave Mechanical Model
of the Atom



Louis de Broglie and Erwin Schrodinger were the next
scientist to explore the structure of the atom. These
scientists suggested that since light behaves like both a
wave and a particle, electrons must do the same.
When Schrodinger performed a mathematical analysis
and considered the electron to be both wave-like and
particle-like, he found that his data worked. His data
worked equally well for all atoms, not just hydrogen.
This model of the atom is called the wave mechanical
model (or the quantum mechanical model).

In the wave mechanical model of the atom
electrons are found in orbitals, which are nothing
like orbits.
 An orbital shows the areas where the electron is
most likely to be found 90% of the time.
Frequently the electron is close to the nucleus,
but occasionally it drifts further away.
 The exact path of the electron is not predictable.
Schrodinger work allowed him to predict the
areas where it most probable to find the electron
in an atom.
11.7 The Hydrogen Orbitals

The probability map for an electron is called an
orbital. The probability of finding an electron
decreases at greater distances from the nucleus,
but it never reaches zero.
 The edge of an orbital is fuzzy, because an
orbital does not have an exact defined size.
Chemists decided to define the area that
contains 90% of the electron probability.

For a hydrogen electron the ground state is the
1s orbital. When the electron absorbs more
energy it moves to a higher state. The higher
energy orbitals have different shapes.
 The hydrogen atom has discrete energy levels,
called principal energy levels. Principal energy
levels are labeled with numbers (like the ground
state is 1). Each principal energy level is divided
into sublevels. Sublevels are notes with letters
(like s, p, d, and f).
 Within
each sublevel are orbitals which
hold electrons.
 The principal energy level (1) contains one
sublevel, or one type of orbital. The only
orbital, the 1s is spherically shaped.
 The
second principal energy level has two
sublevels. These two sublevels are
indicated by the letters s and p.
 The 2s level contains one circular orbital.
The 2p contains three “peanut” shaped
orbitals. Each orbital holds at most two
electrons.
Orbital Labels
 The
number tells the principal energy level
 The letter tells the shape.
 The letter s means spherical orbital; the
letter p means a peanut shaped orbital.
The subscript x, y, or z denotes the
coordinate axes where the peanut lies.
 As
the principal energy level number
increases the distance from the nucleus
increases.
 An orbital should be thought of as a
potential space for an electron.
 Hydrogen may be in its ground state in the
1s orbital, but many other “potential
spaces” exist for the one electron to go
when it is excited.
11.8 The Wave mechanical
Model: Further Development
 The
wave mechanical model of the atom
can be applied to all atoms. The wave
mechanical model of the atom explains the
arrangement of the periodic table.
 The wave mechanical model of the atom
explains the similarities that occur
between elements in the same group,
based on electron arrangements.
 Each
electron has a spin. The electron
can spin in two directions. This spin is
represented with an arrow. Two electrons
in the same orbital must have opposite
spins.
 The Pauli exclusion principal states that an
atomic orbital can hold a maximum of two
electrons, and those electrons must have
opposite spins.
11.9 Electron Arrangements in
the First 18 atoms in the Periodic
Table

Electron fill the orbital with the lowest energy
first; this is called the Aufbau principle.
 The arrangement of electron in an atom can be
represented two ways: electron configuration or
an orbital diagram.
 An electron configuration lists the orbitals that
contain electrons and how many electrons those
orbitals contain.
 An orbital diagram uses a box to represent each
orbital and an arrow to represent each electron.

Example Electron Configuration
• Nitrogen: 1s22s22p3

Example Orbital Diagram:
 When
making orbital diagrams there is
another rule to follow: Hund’s Rule.
 Hund’s Rule states that when more than
one orbital of equal energy is available
electrons go into each orbital before
pairing up.
11.11 Atomic Properties and the
Periodic Table
 Dynamic
Periodic Table
Trends in Atomic Size
 The
atom does not have a sharply defined
boundary that limits its size; therefore the
radius of an atom cannot be determined
directly.
 The atomic radius is one half the distance
between the nuclei of two like atoms in a
diatomic molecule.
 The atomic radius of an atom of an
element indicates the relative size.
Atomic Size in Groups

Atomic size generally increases as you move
down a group.
 As you move down a group electrons are added
to a higher principle energy level and the nuclear
charge increases.
 The enlarging effect of the greater distance from
the nucleus overcomes the shrinking effect of
the increasing charge of the nucleus.
 The outermost orbital is larger as you move
downward, therefore the size of the atom
increases.
Atomic Size in Periods

Atomic size generally decreases as you move
from left to right across a period.
 Each element has one more proton and one
more electron than the preceding element.
 The electrons are added to the same principle
energy level.
 The effect of the increasing nuclear charge on
the outermost electrons is to pull them closer to
the nucleus, therefore the size of the atom
decreases.
Electron Shielding
 The
effect of the increasing nuclear charge
is less pronounced in periods where there
are electrons in energy levels between the
nucleus and the outermost electrons.
 This is because the inner electron shields
the charge of the nucleus. The effect of
shielding is constant within a period
http://images.google.com/imgres?imgurl=http://www.camsoft.co.kr/CrystalMaker/support/tutorials/crystalmaker/resources/VFI_Atomic_Radii_sm.jpg&imgrefurl=http://www.camsoft.co.kr/Cryst
alMaker/support/tutorials/crystalmaker/AtomicRadii.htm&h=450&w=680&sz=80&hl=en&start=3&um=1&usg=__dzl5lBD9l8fzu_rjxzUFr7QtUYo=&tbnid=kjwcl42ApwWlLM:&tbnh=92&tbnw=139
&prev=/images%3Fq%3Dperiodic%2Btable%2Bshowing%2Batomic%2Bradius%26um%3D1%26hl%3Den%26client%3Dfirefox-a%26channel%3Ds%26rls%3Dorg.mozilla:enUS:official%26sa%3DN
Trends in Ionization Energy
 When
an atom gains of loses an electron it
becomes an ion.
 The energy required to overcome the
attraction of the nucleus and remove an
electron from a gaseous atom is called
ionization energy.
Groups and Ionization Energy
 The
energy required to remove the first
electron is called first ionization energy.
 The first ionization energy generally
decreases as you move down a group.
 This happens because the size of the
atom increases and the outermost electron
is further from the nucleus, making it more
easily removed.
Periods and Ionization Energy
 The
first ionization energy increases as
you move from left to right across a period.
 The nuclear charge increases and the
shielding effect is constant, therefore there
is a greater attraction of the nucleus for
the electron and a higher ionization
energy.
Ionization Energy
http://images.google.com/imgres?imgurl=http://www.800mainstreet.com/4/0004-000-IE.GIF&imgrefurl=http://www.800mainstreet.com/4/0004-002Periodic.html&h=340&w=510&sz=39&hl=en&start=19&um=1&usg=__2Ai1eH6TjvM51NHjO4Hyr8FJsIM=&tbnid=Rm_BRv8Igob7pM:&tbnh=87&tbnw=131&prev=/images%3Fq%3Dperiodic%2Btable
%2Bof%2Bionization%2Benergy%26start%3D18%26ndsp%3D18%26um%3D1%26hl%3Den%26client%3Dfirefox-a%26channel%3Ds%26rls%3Dorg.mozilla:en-US:official%26sa%3DN
Ionic Size
 The
atoms of metallic elements have low
ionization energies and thus form positive
ions easily.
 Remember a positive ion is formed when
electrons are lost.
 Atoms of nonmetallic elements however
form negative ions.
 A negative ion is formed when an atom
gains electrons.
Cations (Positive Ions)
 Positive
ions are always smaller than the
neutral atoms from which they form. This
is because the loss of outer-shell electrons
results in an increased attraction by the
nucleus for the remaining electrons.
Anions (Negative Ions)
 Negative
ions are always larger than the
neutral atoms from which they form. This
is a result of the smaller effective nuclear
charge.
Trends in Ionic Size
 Going
from left to right across a row, there
is a gradual decrease in the size of
positive ions.
 Then beginning with group 5, the negative
ions (which are much bigger) gradually
decrease in size as you continue to move
right.
 Going down a group there is a general
increase in ionic size.
Trends in Ionic Size
http://www.chem.umass.edu/~botch/Chem111F04/Chapters/Ch8/IonicRadii.jpg
Electronegativity
 The
electronegativity of an element is the
tendency for the atoms of the element to
attract electrons when they are chemically
combined with atoms of another element.
 Electronegativity is used to predict the
type of bond that occurs between atoms.
Trends in Electronegativity
 Electronegativity
generally decreases as
you move down a group.
 As you move across a period from left to
right the electronegativity increases.
 Fluorine is the most electronegative
element, while cesium is the least
electronegative element.
Trends in Electronegativity
All Periodic Trends
http://www.meta-synthesis.com/webbook/35_pt/general_pt.gif