Chapter 10- Part One
Modern Atomic Theory
Objectives:
•Review history… (10.1)
•Describe electromagnetic radiation (10.2)
•Describe the Bohr atom (10.3)
•Explain energy levels of electrons and
diagram atomic structures for elements
(10.4 & 10.5)
Review…
• Dalton
• Thomson
• Rutherford
– Model doesn’t explain how the negative
electron can stay in orbit and not be
attracted to the positive proton
Electromagnetic Radiation
• Light travels in waves, similar to waves
caused by a moving boat or a pebble
tossed in a pond
• Light is a form of Electromagnetic
Radiation
– Form of energy that exhibits wavelike
behavior as it travels through space
Electromagnetic Radiation
• All waves can be described in 4 ways:
– Amplitude – the height of the wave, results
in the brightness or intensity of the light
– Wavelength (l): distance between
consecutive peaks in a wave
– Frequency (n): number of
waves that pass a given
point in a second
Electromagnetic Radiation
• Speed of light in air: Electromagnetic
radiation moves through a vacuum at
speed of 3.00 x 108 m/s
• Since light moves at constant speed
there is a relationship between
wavelength and frequency:
c = ln
Wavelength and frequency are inversely
proportional
Electromagnetic Spectrum
Quantum Theory
• Wave theory does not explain
– Heated iron gives off heat
• 1st red glow
yellow glow
white glow
– How elements such as barium and strontium
give rise to green and red colors when
heated
Quantum Theory
• Max Planck (1858-1947)
– Proposed that there is a fundamental
restriction on the amounts of energy that an
object emits or absorbs, and he called each
of these pieces of energy a quantum.
• Energy is release in Quanta
Quantum Theory
• A quantum is a finite quantity of energy
that can be gained or lost by an atom
E = hn
E = energy
v = frequency
h = 6.626 x 10-34 J/s
• This constant, h, is the same for all
electromagnetic radiation
Photoelectric Effect
• The emission of electrons by certain
metals when light shines on them
– Albert Einstein (1905) used Planck’s
equation to explain this phenomenon;
• proposed that light consists of quanta of
energy that behave like tiny particles of light
• Photon = individual quantum light (also
known as a particle of radiation)
Photoelectric Effect
• He (Einstein) explained that the
photoelectric effect would not occur if
the frequency and therefore the energy
of each photon is too low to dislodge an
electron.
• Analogy:
– 70 cents placed in soda machine: no soda
– 30 cents more and you will get your soda
Now…
• Light can be described as both particles
and waves
• Dual Wave-Particle Nature of Light was
accepted
• What does this mean for the atom???
Line
Spectrum
• Elements
in gaseous
states give
off colored
light
– High temperature or high voltage
– Always the same
– Each element is unique
• http://home.achilles.net/~jtalbot/data/elements/
Line Spectrum
• Ground state
– Lowest energy level available
• Excited state
– State in which electron has a higher
potential energy than in its ground
state
– Farther from nucleus
– Higher potential energy
Line Spectrum
• Electron falls from higher energy
level to lower one…emits light at a
specific frequency
• Color of light emitted depends on
difference between excited state and
ground state
– See figure 10.5 page 201
Line Spectrum
• Each band of color is produced by
light of a different wavelength
• Each particular wavelength has a
definite frequency and has
definite energy
• Each line must therefore be
produced by emission of photons
with certain energies
Line Spectrum
Line Spectrum
• Whenever an excited electron
drops from such a specific excited
state to its ground state (or lower
excited state) it emits a photon
• The energy of this photon is equal
to the difference in energy between
the initial state and the final
state.
Niels Henrik David Bohr
• 1885-1962
• Physicist
• Worked with
Rutherford
– 1912
• Studying line
spectra
– of hydrogen
Niels Henrik David Bohr
• 1913 – proposed new
atomic structure
– Electrons exist in
specific regions away
from the nucleus
– Electrons revolve
around nucleus like
planets around the sun
The Bohr Atom
• Nucleus with protons and neutrons
• Electrons move in “stationary states” which
are stable (paths or orbits)
• When an electron moves from one state to
another the energy lost or gained is done is
ONLY very specific amounts
• Each line in a spectrum is produced when an
electron moves from one stationary state to
another
The Bohr Atom
•Model didn’t
seem to work
with atoms with
more than one
electron
•Did not explain
chemical
behavior of the
atoms
Wave Matters…
•Louis de Broglie (1924)
•Proposed that electrons might
have a wave-particle nature
•Used observations of normal
wave activity
Wave Matters…
•Erwin Schrodinger (1926)
•Used mathematical
understanding of wave behavior –
devised an equation that treated
electrons moving around nuclei
as waves
•Quantum Theory
Quantum Theory
• Describes mathematically the wave
properties of electrons and other very
small particles
• Applies to all elements (not just H)
Energy Levels of Electrons
• Principal energy levels
– Designated by letter n
– Each level divided into sublevels
• 1st energy level has 1 sublevel
• 2nd energy level has 2 sublevels
• Etc.
Energy Levels of Electrons
Orbitals
• Electrons don’t actually
orbit like planets
• Orbital:
region in space
where there is a high probability of
finding a given electron
– Each orbital sublevel can hold 2 electrons
Orbitals
Each sublevel (orbital) has a specific shape
http://daugerresearch.com/orbitals/
Orbitals
•Pauli exclusion principle: an atomic
orbital can hold a maximum of two
electrons which must have opposite spins
•Electrons can only spin in two
directions
•Shown with arrows
Rules for Orbital Filling
• Pauli’s Exclusion Rule
– No two electrons have the same set of
quantum numbers
• Hund’s Rule
– Electrons will remain unpaired in a given
orbital until all orbitals of the same sublevel
have at least one electron
1s
2s 2p
3s
3p
Rules for Orbital Filling
• Diagonal Rule
– The order of filling
once the d & f
sublevels are being
filled
– Due to energy
levels
Rules for Orbital Filling
Quantum Numbers
• Numbers that specify the properties of
atomic orbitals and their electrons
• Principle Quantum Numbers:
– Symbolized by n, indicates the main energy
levels surrounding a nucleus, which
indicates the distance from the nucleus
(shells or levels)
Quantum Numbers
• Orbital Quantum Number:
– Indicates the shape of an orbital
– (subshell or sublevels)
– s, p, d, f
Principal Quantum #
Orbital Quantum #
1
1s
2
2s, 2p
3
3s, 3p, 3d
4
4s, 4p, 4d, 4f
Quantum Numbers
• Magnetic Quantum Number:
– Indicates the orientation of an orbital
about the nucleus
– Orbital position with respect to the 3dimensional x, y, and z axes
Quantum Numbers
• Spin Quantum Number:
– Indicates two possible states of an electron
in an orbital
Type of Orbital
Number of Orbitals
s
1( )
p
3 (x, y, z) ( ,
,
,)
d
5( ,
,
,
,
)
f
7
Each orbital holds a maximum of 2 electrons
Application of Quantum
Numbers
• Several ways of writing the address or
location of an electron
• Lowest energy levels are filled first
• Electron Configuration: using the diagonal
rule, the principal quantum number (n), and the
sublevel write out the location of all electrons
12C:
32S:
1s22s22p2
1s22s22p63s23p4
Application of Quantum
Numbers
• Orbital filling electron diagram: using
Hund’s rule and the diagonal rule write
out the location of all electrons
• See examples on whiteboard
Homework
•
•
•
•
Worksheet 1
Question #11
Paired Exercises #27-33 odd;
Additional Exercises #54
Chapter 10 – Part Two
The Periodic Table
Objectives:
•Understand the arrangement of the
Periodic Table (10.6)
•Identify connections between electron
configuration and placement on the
periodic table
The Periodic Table
• 1869 – arrangement proposed by Dmitri
Mendeleev
– And Lothar Meyer (different layout)
– Still similar today
– Based on increasing atomic masses and
other characteristics
– Was able to predict properties of elements
not yet discovered….and was correct!
The Periodic Table
• Horizontal rows
– Periods
– Corresponds to outermost energy level
• Vertical Columns
– Groups or families
– Similar properties; reactions
The Periodic Table
• Several systems for naming groups
– Left to right, 1-18
– Roman numerals and A and B
• Used in this book
• Group A: Representative Elements
–
–
–
–
Noble Gases
IA – Alkali Metals
IIA – Alkaline Earth Metals
VIIA - Halogens
• Group B: Transition Elements
The Periodic Table
• Chemical behavior and properties of
elements in a particular family similar
– Have the same outer shell electron
configuration
– Figure 10.15 page 211
• Noble gas configuartion (shortcut)
– Use previous noble gas in square brackets
– Finish with valence electrons
The Periodic Table
• Examples:
– K is 1s22s22p63s23p64s1 or [Ar]4s1
– Ca is 1s22s22p63s23p64s2 or [Ar]4s2
• Write abbreviated configuration for the
following elements:
– Fr
–Y
The Periodic Table
• Arrangement of Periodic Table also
means that elements filling similar
orbitals are grouped
– s block
– p block
– d block
– f block
• Know these blocks…
The Periodic Table Highlights
• The number of the period corresponds
to the highest energy level occupied by
electrons in that period
• The group numbers for the
representative elements are equal to
the total number of valence electrons in
that group
The Periodic Table Highlights
• The elements of a family have the same
outermost electron configuration
– (just different energy levels)
• The elements within each of the s, p, d, and f
blocks are filling the corresponding orbitals
• There are some discrepancies with order of
filling
– (not covered in this book)
Homework
•
•
•
•
•
Worksheet 2
Questions #12-16 even
Paired Exercises #39-44 all; 49-52 all
Additional Exercises #57 & 59
STUDY FOR TEST ON THURSDAY
– Questions answered from 5:00 -5:45 pm
– Test starts at 5:45 pm
– Chapter 11 Lecture will begin at 7:00 pm