Transcript 5 periodicity and atomic structure

Periodicity and Atomic Structure
UU chem 216 chapter 5
Development of the periodic table
• The most important organizing principle in
chemistry (1869). It explained known facts
and made predictions about unknown
phenomena, elements and their properties.
• Periodicity – the repeating of various traits like
atomic radii and valence electrons.
• Mendeleev’s chart lacked the Noble gases,
which were discovered in 1984.
Light and the electromagnetic
spectrum
• Studies of the interaction of radiant energy with
matter have provided immense insight into
atomic and molecular structure.
• Electromagnetic spectrum: chart of all the
different kind of electromagnetic energy from
gamma rays to radio waves.
• Duality of light: in a vacuum (space) travels like
waves with a frequency, wavelength and
amplitude and carries energy as discrete units
called a photon.
Electromagnetic wave
• Wave direction perpendicular to the fields
Calculating a frequency from a
wavelength
• Wavelength X frequency = speed
• Lamda λ (m) v (s-1)
c (m/s)
• The light blue glow given off by mercury
streetlamps has a wavelength of 436 nm.
What is its frequency in hertz?
Electromagnetic energy and atomic
line spectra
• sunlight is white light which is a continuous
distribution of wavelengths of the entire
visible spectrum.
• When shone through a glass prism, the
different wavelengths travel at different rates
and the wavelengths are separated. This is
what happens in rainbow and parhelion
(sundog)
• http://micro.magnet.fsu.edu/primer/java/scie
nceopticsu/newton/index.html
The dispersion of visible light
Atomic line spectra
• Atoms give off light when heated or energetically
excited giving clue to their atomic makeup.
• Light excited contain only a few wavelengths
rather than a full rainbow, giving a series of
discrete lines on dark background – line spectrum
with each element have its own unique spectral
“signature / pattern.” Johann Balmer (series)
visible. Johannes Ryndberg made fit all
• Lyman series – spectral lines in ultraviolet region
• Paschen, Brackett, and Pfund – in infared region
Spectrum pictures
• Wavelength Color 656.2
red 486.1 blue-green
434.0 blue-violet 410.1
violet
Balmer – Rydberg equation
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1/λ =R[1/m2-1/n2] or v=R*c[1/m2-1/n2]
λ - wavelength
R(Ryndberg constant) = 1.097 x 10-2nm-1
M and n represent integers with n>m. If m=1,
then Lyman series, m=2 then Balmer, if m=3
the Paschen series.
Calculation using Balmer – Rydberg
• What are the two longest – wavelength (nm)
in the Lyman series of the hydrogen
spectrum? n=2, n=3.
• What is the shortest – wavelength line (nm) in
the Lyman series of the hydrogen spectrum?
n=infinity
Particlelike properties of
electromagnetic energy
• Albert Einstein - photoelectric effect- irradiating a
clean metal surface with light causes electrons to
be ejected from the metal, with a threshold value
different for each metal.
• The beam of light behaves as it were a stream of
particles (photons) whose energy (E) is related to
their frequency, v (or wavelength λ ) by equation
E=hv or hc/λ
• Higher frequencies and shorter wavelength
correspond to higher energy radiation (lower
frequency and longer wavelength, lower energy)
Photoelectric effect
Quantum
• Intensity of light beam measures the number
of photon, frequency measure of energies.
• Light / matter are quantized – both occur only
in discrete amount, like steps vs a ramp.
• Neils Bohr- his model the energy levels of the
orbits are quantized
Calculating the energy of a photon
from its frequency
• What is the energy (kJ / mole) of photons of
radar waves with a v of 3.35 x 108 Hz
Wavelike properties of matter
• Louis de Broglie (1892-1987) if light can behave in
some respects like matter, then perhaps matter
can behave in some respects like light / particles
• He substituted Einstein E=mc2 into the λ =hc/E to
his equation of λ = h/mv
• The dual wave / particle description of light and
matter is really just a mathematical model. We
can’t see atoms and observe their behavior
directly.
Standing wave review
Calculating the De Broglie wavelength
of a moving object
• What is the de Broglie wavelength (meters) of
a pitched baseball with a mass of 120g and a
speed of 100 mph (44.7 m/s)?
De Broglie electron
• Wavelike pattern for electron
Quantum mechanic and the
Heisenberg Uncertainty principle
• Erwin Schrodinger -1926- proposed the quantum
mechanical model, electron -wavelike not particle
like (orbits)
• Werner Heisenberg – it is impossible to know
exactly where an electron is and what path it
follows. By seeing, energy is imparted to make it
move faster and thus change its position. Not
know both position and velocity. If mass is large
enough, Heisenberg relationship too small to
show problem measuring mass and velocity.
• (delta X) (Delta mv) > h/ 4 pi
Using the Heisenberg Uncertainty
principle
• Assume that you are traveling at a speed of 90
km/h in a small car with a mass of 1250 kg. If
the uncertainty in the velocity of the car is 1%
(delta v = 0.9 km/h) what it the uncertainty (in
meters) in the position of the car? How does
this compare to electron (300pm, size /
diameter is 240 pm)
Wave function and Quantum number
• Schrodinger’s quantum mechanical model is a
differential equation called a wave equation,
since similar to fluid waves. Solution to equation
is called wave functions or orbitals, represented
by Greek psi (Ψ ). Square of wave function is
probability of finding the electron within a
specific region. A wave function is characterized
by three parameters called quantum numbers
represented by letters n, l, m which describe the
energy levels of the orbitals and 3-D shape of the
region of space.
Psi squared
Quantum numbers
• The principal quantum number is (n) is a
positive integer which size and energy primary
dependent, shells around nucleus.
• The angular – momentum quantum number
(l) defines the 3-D shape of the orbital. If N=1,
then l = 0, If N=2, then l= 0 or 1, if N=2, then l=
0,1,2 . Orbitals within a shell is grouped into
subshell. 0=s, 1=p, 2=d, 3=f,
Electron spin and the Pauli Exclusion
Principle
• Magnetic quantum number (ml) defines the spatial
orientation of the orbital in respect to a set or
coordinate axes. For the value of l, there are 2l +2
different spatial orientation.
• If l= 0, then ml= 0, if l=1, the ml = -1, 0, +1; If l=2,
then ml = -2,-1,0,+1,+2
• Spin quantum number (ms)- +1/2 ( )or -1/2 ( )
Electrons closely paired, spin opposite.
Independent of other 3.
• Wolfgang Pauli exclusion principle– no two
electrons can have the same 4 quantum number.
Quantum numbers
The shape of orbitals
• S orbitals – spherical – distance dependent, not
direction, only 1 s subshell per shell. Size increases in
higher shells, Within is a spherical node – zero
probability. No planar node
• P orbitals – dumbbell – shaped, identical lobes on the
either side of a planar node. 3 orbitals at 90* angles
along the x, y, and z axes.
• D orbitals – four clover leaf and 1 dumbbell in a
doughnut. 2 nodal planes
• F orbitals – 7 f orbitals with 8 lobes separated by 3
nodal planes
• http://www.youtube.com/watch?v=sMt5Dcex0kg
Orbital shapes
Quantum mechanics and atomic line
spectra
• When electron absorbs energy from a flame
or electronic discharge they jump to a higher
energy level. It is unstable, so it rapidly
returns to a lower-energy level along with
emission of energy equal to the difference of
the higher to lower level. We observed the
emission of only specific frequencies of
radiation, color wavelengths.
Calculating the energy difference
between two orbitals
• What is the energy difference (kJ / mol)
between the first and second shells of the
hydrogen atom if the lowest-energy mission in
the Lyman series occurs at λ = 121.5 nm?
Orbital energy levels in Multielectron
atom
• Many different interaction for multielectron as
compare to hydrogen with only 1 electron.
• Repulsion of outer-shell electrons by innershell election.
• The nuclear charge felt by an electron if the
effective nuclear charge, Zeff.
• Zeff = Zactual – Electron shielding
• The inner electrons shield the outer electrons
from the full charge of the nucleus.
Electron shielding
Electron configurations of
multielectron atoms
• Predict which orbitals are occupied by
electrons. 3 Rules called the Aufbau (building
up) principle guides the filling order of
orbitals. The resultant lowest energy is called
the ground- state electron configuration.
Several orbitals will have the same energylevel which are said to be degenerate.
Rules of Aufbau principle
• 1. Lower-energy orbitals fill before higher-energy
orbits.
• 2. An orbital can hold only 2 electrons, which
must have the opposite spins. (Pauli exclusion)
• 3. If 2 or more degenerate orbitals are available,
one electron goes into each until all are half – full
(Hund’s rule-far apart due to repulsion
• List n quantum number and the s,p,d,f beginning
with lowest energy and showing the occupancy
of orbital as superscript.
Some anomalous electron
configurations
• Due to the unusual stability of both half-filled
and fully filled subshells.
• Chrominum – predict [Ar] 4s23d4 actually [Ar]
4s13d5
• Copper – predict [Ar] 4s23d9 actually has
[Ar]4s13d10
• Most anomalous elements occur in elements
with atomic numbers greater than z=40,
where energy differences are small.
Common anomalous
configurations
• Element Predicted Electron Configuration
Actual Electron Configuration
• copper, Cu [Ar] 3d9 4s2
[Ar] 3d10 4s1
• silver, Ag [Kr] 4d9 5s2
[Kr] 4d10 5s1
• gold, Au [Xe] 4f14 5d9 6s2
[Xe] 4f14 5d10 6s1
• palladium, Pd [Kr] 4d8 5s2
[Kr] 4d10
• chromium, Cr [Ar] 3d4 4s2
[Ar] 3d5 4s1
• molybdenum, Mo [Kr] 4d4 5s2 [Kr] 4d5 5s1
Electron configurations and the
periodic table
• All element in a group have similar valence-shell
electron configurations.
• Groups 1A,2A form s- block
• Groups 3A-8A form the p-block
• Transitional elements form the d block
• Rare earth (lanthanide /actinide) – form the f
block
• 1s-2s-2p-3s-3p-4s-3d-4p-5s-4d-5p-6s-4f-5d-6p7s-5f-6d-7p
Periodic table and quantum
numbers
Electron configuration arrow diagram
Assigning a ground-state electron
configuration to an atom
• Give the ground-state electrons configuration
of arsenic, z=33, and draw and orbital filling
diagram, indicating the electrons as up or
down arrows.
• Identify the atom with the following groundstate electron configuration
• [Kr] ___ ___ ___ ___ ___ ___ ___ ___ __
Electron configurations and periodic
properties: atomic radii
• Size or atomic radius can be predicted by
electron configuration
• Define an atom’s radius as being half the
distance between the nuclei of two identical
atoms when they are bonded together.