Transcript AP Chem - Unit 2 Chpt7

Chpt 7 - Atomic Structure
•
•
•
•
•
•
Electromagnetic Radiation
Atomic Spectrum - Bohr Model
Quantum Mechanical Model
Orbital Shapes and Energies
Electronic Structure & Periodic Table
Periodic Trends
• HW: Chpt 7 - pg. 329-338, #s 23-27, 37-43, 54, 62, 65, 67-68,
70, 71, 74, 76, 82, 86, 98, 100, 102, 104, 115, 116, 119, 126
Electromagnetic Radiation
 (lamba) = wavelength (m)
 (nu) = frequency (Hertz, Hz or s-1)
E = energy
c = speed of light, 2.9979 x 108 m/s
c =   they are inversely related
Know the relative order of radiation in E,  
1900s Death of Classical Physics
• Black Body Radiation
– Planck’s hypothesis… energy is quantized
E = h or
E = nh n = integer
h = 6.626x10-34 J.s
• Photoelectric effect
– Einstein proposed EM radiation is quantized
– A stream of “particles” called photons
E = h = hc/
• deBroglie  h/mv (wavelength of a particle)
velocity in m/s mass in kg - so units cancel with J
4
Photoelectric Effect
Light with frequency lower than a
specific threshold have no electrons
emitted (no matter how intense it is)
Light with frequency greater than
threshold emits electrons and number
of electrons increases with intensity
Diffraction Pattern in a Crystal
Electron beam is
diffracted off of a
crystal.
Electron exhibits
wave behavior!!!
Davisson Germer experiment - They shared Nobel prize with
GP Thomson which did similar type experiment.
Continuous vs Discrete Spectrum
Continuous spectrum vs. discrete
spectrum (line spectrum)
Absorption vs emission spectrum
Only certain energies are allowed
for the electrons in any atom
Hydrogen Atom
The observed spectrum was explained by Bohr by
proposing the electrons move around the nucleus
in certain allowed circular orbits.
Bohr Energy Expression
• Calculated from hydrogen atom spectrum
E = -2.178x10-18 J (Z2/n2)
Z = atomic number, 1 for hydrogen
n = orbital that the electron is located
• ultimately only good for hydrogen atom spectrum
Quantum Mechanics
• Schrodinger solved the problem mathematically
(no real physical significance) treating electrons
as waves.
H = E 
 is the wave function of the electron’s coordinates in 3
dimensions
• Heisenberg - uncertainty principle
x * (mv) >= h/4
position
momentum
See Heisenberg laser slit video
Orbital shapes and Energies
• Orbitals are simply then a probability distribution of
where the electron could be found.
(left) probability function for s-orbital
(below) Radial probability function for s-orbital
Shapes of p and d orbitals
What can we know about electron?
• 4 Quantum numbers describe the electron in an
orbital.
• n is principle quantum number - relates to size of the orbital,
n = 1, 2, 3, 4,…
• l is angular momentum q.n. - relates to shape of orbital,
l = 0, 1, 2, …, n - 1
•
•
•
•
s-orbital is l = 0
p-orbital is l = 1
d-orbital is l = 2
f-orbital is l = 3
• ml is magnetic q.n. - relates to orientation in space
ml = -l,…,0,…, +l
• ms is electron spin q.n. - relates to spin of electron
ms = - 1/2 or +1/2 (called spin up & spin down or clockwise/counter clockwise)
Quantum numbers
Examples of valid quantum numbers for various orbitals.
In addition, spin +/- 1/2 for each individual orbital.
Energy Levels of orbitals
As we keep adding energy levels,
we see as the principle quantum
number, n, increases the number
of sublevels (types of orbitals)
increases. In addition the energy
spacings get closer together 1s 2s - 3s - 4s - etc. So the energy of
the 4s orbital comes lower than
the 3d. The order need not be
memorized because the elements
in Periodic Table shows it with its
s,p,d,f blocks.
Electron Configuration rules
1. Electron’s occupy lowest energy level first aufbau principle
2. Maximum of 2 electrons in any orbital - Pauli
exclusion principle
–
If 2 electrons occupy the same orbital they have
opposite spins. +1/2 or -1/2 also called spin up
/ down or clockwise / counter-clockwise
3. For degenerate orbitals (the same energy
like the three p, five d, or seven f) use
Hund’s rule, also known as the bus rule only pair up the electrons if necessary.
General s,p,d,f blocks
The periodic table
clearly shows that
after the 3p orbital,
the 4s fills before
the 3d. Likewise, 6s
4f 5d 6p is the order
when the
lanthanides start.
Electron Configurations
A couple of exceptions Cr and Cu groups in the transition
metals promote an s electron to achieve a half-filled and
fully-filled set of d-orbitals because they have more
stability.
Mendeleev’s Original Periodic Table
• Organized by increasing atomic mass and put in
columns by similar properties and reactivities
• Left spaces for undiscovered elements together
with predicted properties - these were confirmed by
experimental results!!!
Periodic Table Trends
some are found in Chpt 8
•
•
•
•
•
Ionization Energy
Electron Affinity
Atomic Radius
Ionic Radius
Electronegativity
Ionization Energies
The ionization energy is the energy necessary to remove an electron
completely from an atom. X --> X-1 + e- The 2nd ionization energy would
remove the next electron, etc.
Notice the trends in this chart a) across the period - general and detailed
b) 1st ion E, 2nd ion E, etc. large jumps associated with core electrons.
1st Ionization Energy Chart
Electron Affinity
EA is the energy change with adding an electron to an atom
X + e- --> X-1
This energy is correlated to thermodynamics, thus atoms that
have a high EA (like to gain e-) the associated E change is
negative (exothermic) the higher EA the more exo it is.
Generally, EA increases up a group and across a period.
Atomic Radius
Radii are estimated from actual
spacing in metals or molecules
Ionic Radius Trends
Ionic radius of most
common ion reported in
picometers.
The size typically
decrease across the
period with a large jump
when going from anion to
cation.
Also, cations are smaller
than their atoms and
anions are larger than
their atoms.
Electronegativity Trends
Electronegativity is the ability of an atom to attract electrons to
itself in a chemical bond. It generally increases across a
period and decreases down a group.
Alkali Metals Periodicity
The alkali metals are shown below with various physical
properties. These are expected trends for other groups
of metals as well.