General Chemistry I
Structure of Atoms
The Wave Nature of Light
to understanding of electronic structure of atoms
The Nature of Waves
for one complete cycle
are length (m)
of cycles for given time (n)
are cycles per second or s-1 (Hz)
of peak from baseline
electric and magnetic components
to each other
at the electric component
electromagnetic radiation moves at the same speed in a vacuum
of light (c)
x 108 m/s
n Wavelength(8) x Frequency (n) = ???
n m x s-1 = m/s
EMR, c = 8n
properties related to wavelength
properly to frequency!
cover a wide range of values!
more appropriate to express wavelengths in units other than meters.
Problems: 1 – 9 odd
Quantized Energy and Photons
can be judged by color of light emitted from an object
hot cooler than white hot
not explain using known laws of physics
energy could only be given off in ‘chunks’
Constant used for energy of single ‘quantum’ of energy
n h = 6.626 x 10-34 J s
n v is the frequency of the emitted
The Photoelectric Effect
energy packets of light called photons.
light a particle or a wave?
‘Dual Nature of Light’
Problems: 11 – 21 odd
Bohr’s Model of the Hydrogen Atom
‘Black Body’ radiation
Hydrogen Line Spectrum
presence of visible lines
The Bohr Atom
observation was pre-nuclear atom.
Rutherford’s discovery, electrons were pictured as small solar system.
physics – electron will spiral into the nucleus.
– used Planck’s idea of quantized energy
certain orbits are ‘permitted.’
n Energy transitions from the Bohr model
n Only these energy transitions are allowed
n Explains why lines are observed.
when n = 4 is 0
loses energy as it approaches the nucleus (n gets smaller).
for electrons in orbits are less than zero!
sign before Rydberg constant
assumed electrons could ‘jump’ orbits.
was energy associated with ‘jump.’
calculate the energy difference
is RH / h?
x 1015 s-1
as Balmer constant!
series is for electrons ‘falling’ to the n=2 level
value means photon is absorbed.
moves to a higher energy level.
sign means photon emitted.
moves to a lower energy level.
n When an electron is promoted it requires energy.
n Fraunhofer lines
n When an electron is ‘demoted’ (returns to a lower
energy state it emits energy
n Line spectra
The Bohr Atom - Limitations
applies to single-atom species
applied to multi-electrons problems arose.
have line spectra, but not where predicted.
Problems: 23 – 29 odd
The Wave Behavior of Matter
light can act like particles, can matter act like waves?
n l =
n h =
Planck’s constant; m = mass (g), v = velocity (m/s)
to uncertainty in position!
mass increases, wavelength decreases
explain why I can’t hit a curve ball!
significant for very small particles
behind electron microscopes
the wavelength, the better the resolution
an electron, l = 0.122 nm
calculate exactly the position and momentum of an object
of the significant wave properties of electrons, cannot determine both the
position and momentum of an electron
the ‘mini solar system’ picture for the atom
Problems: 33 – 37 odd
Quantum Mechanics and Atomic Orbitals
electrons behave as waves, then wave equations should be able to predict
Mechanics or Wave Mechanics
a set of equations to describe electrons based on wave properties
Schrodinger Wave Equation
properties a function of y2
hydrogen, values match what Bohr determined
is we don’t know the location of the electron using the wave function.
is a ‘probability density function.’
yields 3 quantum numbers to describe electrons.
quantum number (n)
to a Bohr ‘orbit.’
quantum number (l)
quantum number (ml)
Assigning Quantum Numbers
have any integral value from 1 to 4.
values range from 0 to n – 1.
l = 0 (s); l
= 1 (p); l = 2 (d); l = 3 (f)
values range from -l to +l.
Quantum Number Principles
n shell will have n subshells
n = 3 there are 3 subshells
l subshell will have 2l + 1 orbitals
l = 2 (d) there are 5
number of orbitals for a shell = n2
4th shell will have 16 orbitals
significant for the periodic table!
n Relative energy levels for the hydrogen atom
n All orbitals in a shell have the same energy
n Explains Bohr atom.
IS TRUE ONLY FOR SINGLE ELECTRON SPECIES!
Representation of Orbitals
n s orbitals
n Shaped like a ball
n Spherical symmetry
n Areas with a low probability of finding an electron
n Nucleus (all)
n Other nodes when n>1.
n p orbitals
n 2 lobes with node at nucleus
n All p orbitals are shaped the same with a different
orientation in space
n Also increase in size as we increase value of n.
lobes for most of the orbitals
sets of orbitals with the lobes between the x/y/z axes
orbital with the 4 lobes on the x/y axes
orbital with 2 lobes (z axis) and a ‘doughnut’!
sets with 8 lobes
sets with 2 lobes and ‘double doughnut’
Problems: 39 – 49 odd
Orbitals in Many-Electron Atoms
descriptions (shapes) for the hydrogen atom orbitals are the same as for other
you have more than one electron, the interaction between the electrons becomes
Effective Nuclear Charge
forces on electrons
to the nucleus
by other electrons
complex to treat in detail, but can make some simplifying assumptions
= Z – S
average # electrons between electron and nucleus
of screening dependent on distance from the nucleus
electrons between the nucleus and electron of interest.
from nucleus increases with n value for a given subshell
from nucleus increases for increasing l
value in a given shell
closer than p closer than d closer than f
decreases as l increases
n For the s (or p, d, f) orbitals
n As n increases, energy increases
n For the 2nd (or 3rd, 4th,…)
n As l increases, energy
n All orbitals are degenerate in a given subshell.
what was thought to be a single line were two closely spaced lines
to ‘electron spin’
quantum number (ms)
of + ½
in hotel room.
Pauli Exclusion Principle
two atoms can have the same set of 4 quantum numbers.
electron has a unique address!
the maximum number of electrons in an orbital must be 2
to understanding why the periodic table works.
Problems: 51 – 61 odd
of electrons in an atom.
will enter lowest available energy level
can hold at most 2 electrons
orbitals will ½ fill with electrons of parallel spin before any fill with
shell (n) and subshell (letter designating the value for l) with superscript indicating the
number of electrons in each.
Li; 1s2 2s1
2 electrons in the 1s subshell (filling it) and 1 electron in the 2s shell (not
n 1s2 2s2 2p6 3s2
n Where does the next electron go (K)?
n We know that 3d is greater than 3s
n We know that 4s is greater than 3s
n Is 4s < 3d???
– [Ar] 4s1
subshell will fill completely (Ca) before the 3d subshell starts to fill (Sc)
can use the periodic table to help predict where the electrons go.
and regions on periodic table defined by their electronic structure!!!
n Electrons in the outermost shell
n For elements in a group, the valence shell
configuration is the same
n Chemical and physical properties are driven by valence
n Families have similar chemical and physical properties
n Inner electrons are ‘core’ electrons
Predicting Electronic Configurations
to use the periodic table to predict these structures
usually occur when atom can have a more stable arrangement using ½ filled or completely
n Ag –
predict [Kr] 5s2 4d9
n Ag –
actual [Kr] 5s1 4d10
n Ag – [Kr] 5s1 4d10
n Energy levels are dynamic, not static
n Above electronic structure has only ½ filled and
completely filled subshells.
n HOW THE &%## AM I SUPPOSED TO KNOW THAT!
n I expect you to be able to predict using the periodic
table (Ag would be [Kr] 5s2 4d9)
n I expect you to be able to explain why an anomaly
would occur GIVEN THE CORRECT CONFIGURATION!
Problems: 63 – 69 odd